LIBRARY OF CONGRESS. 

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UNITED STATES OF AMERICA. 



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INORGANIC CHEMISTRY 



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.mm 



Inorganic Chemistry 



FOR 



BEGINNERS 




BY 



Sir henry ,R0SC0E, F.R.S., D.C.L., LL.D., M.P. 



ASSISTED BY 



JOSEPH LUNT, B.Sc. (Vict.), F.C.S. 



With one hundred and eight Illustrations in the Text 



MACMILLAN AN 

AND LONDON 

1893 




D Q0.^9lZ^J 



All rights reserved 



■V. . 



Copyright, 1893, 
By MACMILLAN AND CO. 



Xorfajooti Prrss : 
^ J. S. Cushir 



y^ J. S. Gushing & Co. — Berwick & Smith. 

- ^ Boston, Mass., U.S.A. 



«^^ 



PREFACE 

The publishers of the Elementary Lessons in Chemistry 
have brought before my notice the want of a work for 
those beginning the study of the science, in which the 
elementary principles of chemistry are more fully treated 
than is the case in the Lessons, whilst the description of 
the elements and their compounds is restricted to a few 
well-chosen typical examples. In the present pages I 
have endeavoured to fulfil the above requirements, and 
I have found no selection more suitable to the needs of 
beginners than that made many years ago for this purpose, 
and still found to be effective, by the Department of Science 
and Art. It will be seen that whilst the ground covered 
in this Httle book is confined to the discussion of a portion 
only of the non-metallic elements, a greater amount of 
detail is in each case given than was found to be possible 

in the Lessons. 

H. E. R. 

August 1893. 



CONTENTS 

PART I 

ELEMENTARY PRINCIPLES 

LESSON PAGE 

1. Solids, Liquids, and Gases i 

2. The Air — Introductory 8 

3. Water, Introductory — Mechanical Mixture and 

Chemical Combination — Indestructibility of 
Matter 14 

4. Elements and Compounds — Symbols and Formul/e — 

Distribution of the Elements' 21 

5. Combination in Definite and Multiple Proportions 

— Dalton's Atomic Theory— Atomic and Molecu- 
lar Weights — Calculations 27 

6. Physical Measurements. — Standards of Length, Volume, 

and Weight used in Chemical Experiments — Methods of 
Weighing and Measuring — The Metric System — The Ther- 
mometer — Measurements of Temperatures — Conversion of 
Thermometric Scales 38 

7. Physical Properties of Gases. — Relation of Volume to 

Temperature and Pressure — Dalton's Law — Boyle's Law — 
Calculation of Volumes from Weights — Reduction to Nor- 
mal Temperature and Pressure [NTP] 50 

8. Physical Properties of Gases — Continued. — Relations 

between the Combining Volumes of Gases — Avogadro's Law 

— Diffusion, Liquefaction, and Solidification of Gases . . 59 

vii 



CONTENTS 



PART II 

SYSTEMATIC STUDY OF CERTAIN NON-METALLIC ELE- 
MENTS, AND THEIR MORE IMPORTANT COMPOUNDS 

LESSON PAGE 

9. Hydrogen 67 

10. Oxygen and the Oxides, Hydroxides, Acids, Bases, and 

Salts 74 

11. Ozone. — Preparation — Properties and Composition . . . 85 

12. Compounds of Hydrogen and Oxygen H.2O and H.2O.2. 

— Water (Hydrogen Monoxide) — Determination of its 
Chemical Composition by Eudiometric and Gravimetric 
Synthesis, and by Electrolysis — Hydrogen Dioxide . . 91 

13. Heat Relations of Water. — Expansion and Contraction — 

Point of Maximum Density — Tension of Vapour — Evapo- 
ration — Melting and Boiling Points — Latent Heat — Freez- 
ing Machines — Specific Heat — Dulong and Petit's Law . 102 

14. Water as a Solvent — Water of Crystallisation, 

Efflorescence, Deliquescence — Solubility of Gases 

— Natural Waters — Temporary and Permanent 
Hardness and the Softening of Water — Distilla- 
tion and Purification 116 

15. Nitrogen and Air. — Pressure, Temperature, Humidity, and 

Extent of the Atmosphere — The Barometer — Chemical 
Composition and Analysis of Air — Action of Animals and 
Plants on the Air — Ventilation 129 

16. Compounds of Nitrogen and Oxygen, N2O, N2O2, N2O3, 

N2O4, AND N2O5, Nitrous Acid, HNO2, Nitric Acid, 
HNO3, THE Nitrites and Nitrates . . . .149 

17. Compounds of Nitrogen and Hydrogen, Ammonia 

(NH3), AND THE Ammonium (NH4), Compounds . . 162 

18. Chlorine, Hydrochloric Acid, and the Chlorides . 170 

19. Sulphur, Sulphuretted Hy^drogen, and the Sulphides 187 

20. Oxides and Oxy-Acids of Sulphur. — Sulphur Dioxide, 

Sulphurous Acid, and the Sulphites — Sulphur Trioxide, 
Sulphuric Acid, and the Sulphates 200 

21. Carbon and its Allotropic Modifications — Carbon 

Monoxide — Carbon Dioxide — Methane — Acety- 
lene — Ethylene — Coal Gas and Flame . . . 216 



CONTENTS 



APPENDIX 

COMPARISON OF THE METRICAL WITH THE COMMON 
MEASURES 

PAGE 

Measures of Length 241 

Measures of Surface 241 

Measures of Capacity . . , 242 

Measures of Weight 242 

INDEX 243 



PART I 

ELEMENTARY PRINCIPLES 
LESSON I 

SOLIDS, LIQUIDS, AND GASES 

The existence of different kinds of solids, such as wood and 
iron, and of different kinds of liquids, such as oil and water, has 
been recognised by men from the very earliest times ; but that 
similar differences exist amongst invisible gases has only been 
ascertained within a comparatively recent period. These differ- 
ences can, however, be readily shown by experiment. 

ExPT. I . Behaviour of different Gases towards a burn- 
ing Taper. — Here, for example, are three bottles filled with 
colourless invisible gases. I remove the stopper and plunge 
a lighted taper into the first one, and we obser^'e that it 
is at once extinguished. I put the burning taper into the 
second vessel, when we notice that it burns with very much 
increased brilliancy'; and if I bring it to the mouth of the third 
vessel, we find that the gas takes fire and burns with a pale blue 
flame. 

ExPT. 2. Heavy and Light Gases. — Not only do these 
three gases differ essentially from one another, in their behaviour 
towards a burning taper, as we have seen, but they also differ 
greatly in weight. Thus, if we take three bottles of the same 

I 



HEAVY AND LIGHT GASES 



PART I 



three gases, and in the same order, we can easily prove that 
the first not only puts out a taper, but is* a very heavy gas 
and may be poured downwards from one vessel to another 
like water. Although we cannot see the gas pour from one 
vessel to another, it is easy to show that it has passed from B to 
A (Fig. i), by means of a lighted taper, for the flame is put 
out in A, into which the heavy gas has been poured, w^hilst it 
burns quite well in B, which at first contained the heavy gas, 
but which now contains ordinary air which enters as the heavy 
gas leaves. In the same way we can show that the third vessel 
contains a very light gas. This time we shall be able to pour 
it upwa7'ds (see Fig. 2). Having inverted the gas jar B, I take 




A/rK 




Fig 



off the cover and apply its mouth to the mouth of A, also 
inverted, and gradually lower the upper end of B, so as to pour 
the gas upwards into A. As the light gas ascends into A, 
the air is driven out, and as it leaves B, its place is taken by 
ordinary air. This is easily seen to be the case by applying a 
lighted taper to the mouth of both jars, A is now seen to contain 
the inflammable gas, because it takes fire, whilst in B, the taper 
burns just as in ordinary air. 

ExPT. 3. Heavy and Light Gases in the Balance. — 
Another method of proving that there are heavy and light 
gases may be shown as follow^s. At the two ends of the 
arms of a balance are fixed two thin beaker-glasses (Fig. 3), 
one inverted and the other not. These are exactly balanced 



LESSON 



HEAVY AND LIGHT GASES 



when both are filled with air, by adding a little sand to 
the lighter end. If now we have three more bottles of the 
different gases, and pour the contents of the first one downwards 
into beaker A, we see that the heavy gas weighs it down and 
the beaker descends although nothing can be seen to have 
entered the beaker. In the same way if w^e again equipoise the 
vessels, the contents of our third jar, poured upwards into B 
(Fig. 4), make it ascend, showing it to be filled with a very 




Fig. 3. 

light gas. Whilst our second gas, if we empty the beakers of 
the first gases and put them again in equilibrium, will be found 
neither to weigh down A nor raise B, because it is just about as 
heavy as ordinary air. 

ExPT. 4. Behaviour of Gases towards Lime-Water. 
— Again, if we take some clear lime-water* and pour it into jars 
of the three gases used in the first experiment, we find that 
in the first one it becomes quite white and milky, whilst in the 
two others it does not do so. This is because the first gas 
* See footnote to p. 139. 



THE THREE STATES OF MATTER 



PART I 



forms with the Ume in the lime-water a white substance ahnost 
exactly like chalk, which floats about in the water and turns 
it milky. The other two gases do not form this white substance 
with lime-water and produce no change in it. 

These, then, are simple examples of the experimental method 
which the chemist uses, and it is by experimenting that the 
chemist has been able to find out all that he knows about 




Fig. 4. 



the different materials of which Earth, Air, and Water are 
made up. 

ExPT. 5. Substances can Exist in any of the three 
States. Water. — We now see that every substance may be 
classified as either solid, liquid, or gas ; but it will be evident, 
after a little consideration, that many substances can exist in 
more than one of these forms. Let us take the familiar 
example of water. If we place a saucer full of water out in the 
air in winter when it is freezing, we see that the water is turned 
into solid ice. Yet if this be brought into a warm room the ice 



SOLIDS. LIQUIDS, AND GASES 



melts, and we again get liquid water, whilst if we pour this into 
the kettle and put it on the fire, we soon get the water in the 
state of gas or steam which, just as it rushes out of the spout 
of the kettle is quite invisible, but a little way from the spout 
is seen as a cloud or mist consisting of very small drops of 
liquid water. 

ExPT. 6. Gaseous Water or Steam invisible. — That 
the true steam (or gaseous water) is invisible, is easily shown 
by boiling the water in a 
glass flask fitted with a cork 
and bent tube, as in Fig. 5. 
When the water boils, the 
cloud of condensed steam is W: 
seen to rush out from the ^ 
little glass tube, but inside 
the flask, above the boiling 
water, nothing can be seen, 
but still the space must be 
full of gaseous water, for 
directly it gets into the cold 
air, it becomes condensed 
into a cloud of minute par- 
ticles of liquid water. Water, 
then, exists in the solid state, 
as ice, snow, and hail ; in 
the liquid state as water; 
and in the gaseous state as 
steam. 

ExPT. 7. Sulphur as a 
Solid, Liquid, or Gas. — 
In the same way Sulphur, a 
yellow solid body, if heated in a glass retort, will first melt to 
a liquid, and if this be heated still further, it will give off a 
very dark-red gas which condenses in the cold neck of the 
retort to liquid, and then to solid sulphur.* 

ExPT. 8. Iodine in the Gaseous State. — Similarly the 
black metallic-looking substance Iodine, if dropped into a hot 
flask, immediately fills it with a beautiful violet gas, which, as 




^ Fig. 94 shows the method of heating a substance in a retort, 



6 SOLIDS, LIQUIDS. AND GASES part I 

it cools, condenses on the sides of the flask as a sparkling mass 
of black crystals. 

Even the liquid metal mercury can be made to assume the 
gaseous state by strongly heating it in a glass retort. The 
gaseous metal passes into the neck of the retort and is there 
cooled and condensed to shining little drops of quicksilver, 
which run into the small flask placed to receive them. If 
these be put in a freezing mixture which produces iritense 
cold, the liquid metal becomes solid and can be hammered 
or moulded into various shapes like a piece of lead. Most 
solid substances, if heated strongly enough, unless they de- 
compose, will melt to a liquid, and liquids, if heated strongly 
enough, will boil and become gaseous. Similarly, all gases 
will condense to liquids, and all liquids, if they are made cold 
enough, will assume the solid state. Some gases are easily 
condensed to liquids, such as steam, whilst others require to be 
cooled to a point far colder than ice or even than the coldest 
Arctic winter, and moreover, must at the same time be sub- 
jected to very great pressure, before they condense to liquids. 
Similarly with liquids, some are easily solidified, such as water, 
whilst others must be cooled down to a very low point before 
they solidify. We must, however, remember that in all these 
transformations of solid into liquid, and of liquid into gas, and 
vice versa, the essential nature or chemical composittoji of the 
substances is not altered, they are va^rtly physical changes. 



What we have learnt 

In our first Lesson we have learnt that not only are there different kinds 
of solids and liquids, but also different kinds of gases, which, although 
quite colourless and invisible, may be shown to be different by means of 
suitable experiments, such as that with a burning taper or by that of pour- 
ing them, from one vessel to another, or of weighing them on a balance. 
The three gases considered are (i) a very heavy gas, which puts out the 
flame of a burning taper and turns lime-water milky. This is called Carbonic 
Acid Gas. (2) A very light gas which burns in air with a blue flame and 
does not turn lime-water milky. This is called Hydrogen. (3) A gas just 
about as heavy as ordinary air, in which the taper burns with greatly in- 
creased brilliancy, and which does not turn lime-water milky. This is called 
Oxygen. We have also learnt that many substances, such as water, sul- 
phur, iodine, and mercury, can exist in any of the three states of matter; 



SUMMARY AND EXERCISES 



e,g., water, as solid ice, snow, and hail ; as liquid water, and as gaseous 
steam. Similarly, gases may be condensed to liquids and liquids to solids 
without altering their chemical composition. 



Exercises on Lesson I 

1. How can you distinguish, oxygen from hydrogen, and carbonic acid 
from either ? 

2. Mention several solids, liquids, and gases. 

3. How can you show that gaseous water is invisible ? 

4. What happens when iodine is heated in a glass flask ? 

5. Describe any experiments which show that there are different kinds 
of colourless invisible gases. 

6. What happens when lime-water is poured into (i*) oxygen, (2) hydro- 
gen, (3) carbonic acid gas ? 

7. What is meant when we say that water can exist in any of the three 
states of matter ? 

8. What happens when a burning taper is placed in (i) oxygen, (2) 
hydrogen, (3) carbonic acid ? 

9. How can you show by experiment that hydrogen is a light gas and 
carbonic acid gas a heavy gas ? 



LESSON II 



THE AIR, INTRODUCTORY 



Let us now try to find out something about the very com- 
monest gas with which we are acquainted, namely, the 
atmosphere or air whicli surrounds us. 

ExPT. 9. Burning Phosphorus in Air; the Air con- 
tains two Gases. — Let us take a large glass bell-jar A, 

furnished with a stopper at its neck, 
and place it over water in a glass 
dish B, so as to enclose air inside 
the bell-jar (Fig. 6), the water rising 
to the level of the first mark, so as 
to enclose five volumes of air. We 
will float on the water a small porce- 
lain dish D in which a small piece 
of dry phosphorus* has been placed. 
Taking out the stopper of the jar 
and introducing a glass rod, the 
end of w^hich has been heated in 
a gas flame, I will touch the small 
piece of phosphorus with it, and 
this being very inflammable im- 
mediately takes fire and burns with 
a bright white light. Quickly re- 
placing the stopper we see that the 
jar becomes filled with dense white 
fumes, and that after a time the phosphorus begins to burn dimly 
and soon afterwards the flame dies out. Then the dense white 

* Phosphorus must always be cut under water, and must not be handled 
by warm fingers. 




Fig. 6. 



THE AIR, I INTRODUCTORY 



fumes begin to dissolve in the water and disappear. When 
the gas inside has cooled down to the ordinary temperature of 
the air, I pour water into the glass trough until the water 
both inside and outside the bell-jar is at the same level (Fig. 6). 
It will now be seen that we have less gas in our bell-jar than 
when we started the experiment ; and if we exactly measure the 
amount or volume of air we begin with, we shall find that we 
have just about i of the original volume left, for the water has 
now risen to the second mark on the bell-jar, showing that the 
original five volumes have become four volumes. 

Now let us see how the gas remaining will behave towards 
a burning taper. On removing the stopper and plunging in 
the lighted taper the flame will at once be extinguished. 
Evidently, the gas remaining in the bell-jar after phosphorus 
has been burnt in it is something different from ordinary air, 
because in ordinary air a taper will burn, but in this gas it will 
not. (This gas is called Nitrogen.) Perhaps that may have 
been the reason why the light of the phosphorus was put out. 
Let us see if that is so. On taking out the stopper and remov- 
ing the bell-jar from the glass trough we shall find, if we have 
taken sufficient phosphorus, that it again bursts into flame in 
ordinary air if touched with a hot glass rod, showing that the 
flame of the phosphorus did not go out because all the phos- 
phorus had been burnt away and used up, but rather because 
something in the air which allows things to burn has been 
used up, and ^ the volume of the air consists of this gas which 
is used up when phosphorus burns in it. 

ExPT. lo. Lavoisier's Experiments on Air. — We will 
now try to find out something more about the air with which 
we are surrounded. We know already that | of it is com- 
posed of nitrogen w^hich does not allow^ things to burn in 
it. What is the other I composed of? To find this out, 
I will describe an experiment which was made a long time 
ago by the French chemist Lavoisier. Into a glass bal- 
loon or retort A, having a long straight neck Lavoisier brought 
a few ounces of dry mercury or quicksilver ; he then bent 
the neck so that when the balloon rested over the flame B, 
the end of the neck appeared above the surface of the mercury 
contained in the large glass dish C (Fig. 7). Thus the air 
enclosed in the bell-jar D was in communication with that 



LAVOISIER'S EXPERIMENTS ON 



PART I 



in the retort. The volume of air contained in the bell-jar and 
retort was first measured, and then the mercury in the retort 
was made very hot by a flame placed underneath. For the 
first few hours no change occurred, but then red coloured specks 
and scales began to make their appearance. These increased 
in number for some time, but after a time no further for- 
mation of this red substance was observed ; so after the 
experiment had been continued for twelve days the fire was 
removed and the volume of tlie air was seen to have under- 




Fig. 7. 

gone a remarkable diminution, which when measured was 
found, just as in our experiment with the phosphorus, to 
be about ^ of the original volume. Where has this ^ of 
the air gone to? It has burnt up a small portion of the 
mercury in the retort and formed the red powder on its 
surface, and this red powder contains the part of the air 
(\ of its original volume) which is concerned in burning 
(either a taper or a fire or anything else). Lavoisier next 
very carefully collected all the red powder and placed it 



LESSON 



THE COMPOSITION OF AIR 



in a long tube (Fig. 8), connected with a graduated glass cylin- 
der. On heating the red powder very strongly he found that it 
gave off exactly the same volume of gas which had disappeared 
in the previous experiment, and on testing it, found it to be the 
gas in which things burn so readily (this is called Oxygen).* 
This gas is, therefore, termed a supporter of combicstion. That 
oxygen does support combustion most powerfully is shown by 
introducing a chip of wood with only the least spark on it into 




Fig. 8. 



the gas, when the wood will immediately burst into bright flame. 
This is a test for oxygen gas. 

Let us now consider the meaning of these experiments a little 
more closely. 

In the first experiment (Expt. 9) we burnt phosphorus in air, 
and we now know that the air contains oxygen and nitrogen. 
When phosphorus burns, it is cojubining chemically with oxygen, 

* Priestley had previously prepared oxygen by heating mercuric oxide 
(seep. 74). 



12 COMBINATION AND DECOMPOSITION PART i 

and during this chemical action heat and light are evolved. In 
very many cases of chemical action we shall find that heat and 
sometimes light is given oiT. Here we have a good example of 
chemical coinbinatioji ; the yellowish wax-like solid combines 
with the oxygen gas to form the white solid called oxide of phos- 
phorus, the fumes of which filled our bell-jar. We notice here 
that in this chemical combination two diiTerent substances have 
united chemically to form a third substance differing entirely in 
its properties from the materials of which it is made up. We 
next notice that the fumes of the white solid oxide of phosphorus 
soon disappear, and that is because another chemical action 
takes place. The oxide of phosphorus unites chemically with 
the water in the glass dish, and forms a substance which dis- 
solves in the water, this is called phosphoric acid. That the 
water in the trough contains a new substance may be shown by 
pouring into it a little blue litmus solution, which immediately 
turns red, proving that an acid is present. This is a lesl for an 
acid. Thus we have : — 

Phosphorus + Oxygen = Oxide of phosphorus. 
Oxide of phosphorus -f- Water = Phosphoric acid.* 

The nitrogen present in the air takes no part in these chemical 
reactions. 

Now let us further consider the results of Expt. lo. In the 
first part we have another example of chemical combination. 
The silvery liquid mercury, or quicksilver, unites chemically with 
the oxygen of the air and forms a red powder called oxide of 
mercury. In the second part, the oxide of mercury is chemically 
decomposed, yielding mercury and oxygen. This is an example 
of chemical decomposition^ and we notice that one substance in 
decomposing yields two entirely different substances. The pow- 
der is a red solid, and its products of decomposition are a sil- 
very liquid and a colourless gas. Thus we have : — 

(Combination) Mercury -\- Oxygen = Oxide of mercury. 
(Decomposition) Oxide of mercury == Mercury + Oxygen. 

* The sign + is here used to signify " and," or " together with," whilst 
the sign = means " yields." We shall afterwards find that these signs may 
mean more than I now need explain. 



LESSON II SUMMARY AND EXERCISES 13 

In this case we notice that mercury and oxygen must be raised 
to a high temperature before they will combine whilst at a still 
higher temperature the oxide of mercury decomposes. 

Meaning of Analysis and Synthesis. — When mercury 
and oxygen combine together, a synthesis (joining together) of 
mercuric oxide from its elements is said to take place. When 
mercuric oxide is decomposed by heat it is said to be an mialysis 
(or splitting asunder) of that compound. 



What we have learnt 

In our second Lesson we have burnt phosphorus in air enclosed in a 
bell-jar, and noticed that the volume of air is diminished by \ leaving | of a 
gas called nitrogen which puts out a burning taper (but which will not 
turn hme-water milky, distinguishing it from carbonic acid). During the 
burning, the phosphorus is combining chemically with the oxygen of the 
air, forming oxide of phosphorus which combines with the water to form 
phosphoric acid, a substance which turns blue litmus solution red. 
Lavoisier's experiment shows us that the \ of the original air which 
disappears, combines with mercury to form a red solid, called oxide of 
mercury, this, on heating strongly, gives off the oxygen which it contains, 
and this gas is seen to occupy a volume equal to that which was lost in the 
first instance. The gas oxygen is in this manner separated from the 
nitrogen, and we find that air is made up principally of nitrogen (four 
volumes) and oxygen (one volume). 

We have had four examples of chemical action, three of chemical 
combination, viz. (i) phosphorus and oxygen, (2) oxide of phosphorus 
and water, (3) mercury and oxygen; and one example of chemical 
decomposition, viz. oxide of mercury into mercury and oxygen. 



Exercises on Lesson II 

1. Describe an experiment showing that air contains t its volume of 
nitrogen ; state some properties of nitrogen. 

2. How can you separate the oxygen from the nitrogen of the air ? 
How would you test for oxygen ? 

3. What evidence have you that an acid is formed when you burn 
phosphorus in air confined over water ? How would you test for an acid ? 

4. Give some examples of chemical combination and decomposition. 

5. What is meant by the terms analysis and synthesis ? 



LESSON III 

WATER, INTRODUCTORY — MECHANICAL MIXTURE AND 
CHEMICAL COMBINATION — INDESTRUCTIBILITY OF 
MATTER 



The ancients recognised only four (so-called) elements — 
Earth, Air, Fire, and Water. We have already seen that the air 
consists principally of two gaseous substances with widely dif- 
ferent properties. Now let us examine anotljer of the so-called 
elements of the ancients, viz. water. Here again we will make 




Fig. 9. 

use of the experimental method of examination which chem- 
ists use. 

ExPT. II. Action of Sodium on Water. — Here is a small 
bottle containing the soft white metal sodium, which must al- 
ways be kept under mineral oil because it is rapidly acted upon 
by the moisture in the air. If I take a lump of the sodium and 

14 



LESSON III 



ELECTROLYSIS OF WATER 



15 



cut off a small piece with a knife (for sodium, although a 
metal, is very soft and can easily be cut), and then introduce 
the little piece of sodium into a small wire cage furnished 
with a handle, and plunge it into a glass dish full of water 
over which is inverted a gas cylinder also filled with water 
(Fig. 9) ; we notice at once a violent action going on and 
bubbles of gas rise and fill the gas cylinder placed to receive 
them. Here the sodium has decomposed the water and driven 
out or liberated one of its components (which we can recognise 
as hydrogen if we apply a lighted taper to the gas) ; whilst 
the sodium has united or combined with the remaining por- 
tion of the water to form a new substance, caustic soda, 
which remains dissolved. That a new substance is dissolved 
in the water is easily shown by adding some red litmus solu- 
tion, it is immediately turned blue. This is a test for alkalis, 
and caustic soda is an alkali. 



Sodium + Water = Caustic soda + Hydrogen. 

Evidently then w^ater contains hydrogen. What else does it 
contain? A further experiment will tell us. 

ExPT. 12. Electrolysis of Water into Oxygen and 
Hydrogen. — Here is a bat- 
tery * which will furnish a cur- 
rent of electricity ; this has the 
power, if passed through vari- 
ous bodies, of decomposing 
them into their components. 
In Fig. 10 we see the appa- 
ratus which is used for passing 
a current of electricity through 
water. The two wires from 
the battery terminate in plati- 
num plates, and over each one 
hangs a little glass cylinder 
filled with acidulated water. 
Directly we allow the current 
to pass through, we notice that 
from each platinum plate small bubbles of gas begin to arise, 

* See Fig. 48, p. 95. 




i6 MECHANICAL MIXTURE AND part I 

and these pass up into the two tubes placed to receive them, 
and soon we observe that one tube is filling with gas much 
faster than the other, and at last, when one of the tubes is quite 
full the other is only half full. The current of electricity, 
in passing through the water, has rent its components asunder 
and separated them one at one pole or electrode * and the other 
at the opposite one. Let us now ascertain what these gases are. 
For this purpose we will take the tube, which is full, out of the 
water, keeping it closed with the thumb all the time, and then 
try to light the gas with a taper. We see, as in the last experi- 
ment, that the gas takes fire and burns with a blue flame, so this 
is hydrogen. If now we put a glowing chip of wood into the 
other tube half filled with gas, it immediately bursts into flame 
showing that the gas is oxygen. This experiment teaches us 
that from water we can obtain two volumes of hydrogen and one 
volume of oxygen. 

Water = Hydrogen + Oxygen. 

The operation of deco7nposing a substance by passing through 
it a current of electricity is called electrolysis. 

Let us try to understand more clearly what is meant by 
chemical combination, and how to distinguish this from a mere 
mechanical mixture of two bodies. 

ExPT. 13. Sulphur and Iron mechanically mixed and 
chemically combined. — A simple experiment will illustrate 
what this means. I will take a small quantity of fine iron 
filings, place them in a mortar, and add to them about the 
same quantity of yellow flour of sulphur ; on rubbing these two 
together a greenish powder is produced. Examined by the 
naked eye the separate particles of iron and sulphur cannot be 
seen ; if, however, we use a magnifying glass, they can be dis- 
tinguished lying side by side, and by using a small magnet, 
the iron can be drawn out and the sulphur left behind. Next I 
place some of this greenish pow^der in a test-tube, and heat it 
over a gas flame. We notice that the mass becomes red-hot. 
Having allowed it to cool I break the end of the tube in a clean 

* The poles or electrodes are the platinum plates in which the two wires 
from the battery terminate. 



LESSON III CHEMICAL COMBINATION 17 

mortar ; a black solid mass (sulphide of iron) is seen to have 
been formed in which the particles of iron and the particles of 
sulphur cannot be noticed to lie side by side as before, even 
when viewed by a powerful microscope, nor can I draw out any 
particles of iron by means of the magnet. 

In the first part of the experiment I have made a mechanical 
mixture of the two powders. In the second I have brought 
about a chemical combination. By no mechanical process can 
the sulphur and the iron then be separated. This, howxver, 
can be done by chemical means ; and if I could weigh the 
sulphur and the iron to begin with, and afterwards weigh 
the black mass, I should find that the weight of the latter is 
exactly equal to the weight of the original materials which I 
used. 

Matter is Indestructible. — In all the different changes 
of this sort which chemists have been able to bring about, 
. they have found that if care be taken, the weight of the chemical 
substances produced is always equal to the weight of the 
substance or substances before the change. In other w-ords 
matter is indestructible^ and this is true even where an apparent 
loss of matter takes place. Thus, when a candle burns, the 
wax of the candle gradually disappears, and of course the weight 
of the candle as gradually diminishes. A simple experiment, 
however, show^s that the materials of the wax are not lost or 
destroyed but have passed into a state in which we cannot see 
them, but their existence can easily be shown by the following 
experiment. 

ExPT. 14. Carbonic Acid Gas produced when a Candle 
burns. — I place a burning candle for a few minutes in a clean 
bottle filled with air. On withdrawing the candle I pour into 
the bottle some clear lime-water, and notice that this at once 
becomes milky, which it does not do if I pour some into another 
bottle filled with air, in w^hich a candle has not been burnt ; 
proving that the invisible gas, carbonic acid, which turns lime- 
water milky has been formed by the burning of the candle. 
Some of the wax has been changed, by burning, into carbonic 
acid gas, a compound of carbon, one of the materials of the 
wax, with the oxygen of the air. It is easy to show that the 
white wax contains carbon, for some carbon goes away as 
soot or smoke w^hen a candle is burnt, and this we notice .by 



i8 INDESTRUCTIBILITY OF MATTER part I 

pressing a sheet of white paper quickly on to the flame so as 
not to burn it, when the paper becomes blackened with soot or 
carbon. 

ExPT. 15. Water (gaseous) also produced when a 
Candle burns. — But besides carbonic acid, water is formed 
when a candle burns. This is shown by holding a cold dry 
bright glass, such as a tumbler, over the flame of the taper. 
The bright glass becomes at once dim, and small drops of water 
are seen like little drops of dew inside the glass. This is 
because a candle not only contains carbon but also hydrogen, 
and this in burning or combining with the oxygen of the air 
forms water. For just as we were able to obtain oxygen and 
hydrogen by decomposing water, so by combining oxygen and 
hydrogen we again obtain water. This water is given off from 
the candle in the invisible gaseous state as steam, but in coming 
in contact with the cold glass it is cooled sufficiently to form 
liquid water. If, then, what I have said is true, namely that 
when a candle burns a chemical combination takes place be- 
tween the materials of the candle and oxygen gas, if we could 
collect the whole of the carbonic acid and the water formed 
when the candle burns, the weight of these products must be 
actually greater than the weight of the candle which has dis- 
appeared, because oxygen, like every other gas, possesses weight, 
and the increase of weight which is observed must be exactly 
equal to the weight of the oxygen which has combined with the 
carbon and hydrogen of the candle. 

ExPT. 16. No loss of Weight, but an Increase, ob- 
served when a Candle burns. — First, we will take a wide 
tube (A) 5-inch wide and 10 inches long, and into the bottom 
fit a cork which is perforated with holes to admit air, and then 
fix our candle inside the tube, and resting on the perforated 
cork. Then to the top of the tube we will fit another cork, 
through which a piece of glass tubing passes to a (J-shaped 
tube, which is filled with small pieces of a white solid substance 
called caustic soda, which has the power of combining with 
and, therefore, absorbing carbonic acid and water. Next, 
having placed our unlighted candle in position, we will carefully 
weigh the whole arrangement as shown in Fig. 11 on one arm 
of a pair of scales. 

The tube C is now connected by an india-rubber tube to a 



LESSON III 



EXPERIMENT WITH CANDLE 



19 



perforated cork at the top of a large bottle of water, which has 
a tap at the bottom for drawing off the water. If we allow 
the water to run out of our large bottle, air must pass through 
the perforated cork through the U tube, and fill the upper part of 
the bottle as the water flows out ; whilst, if we light the candle, 
not only is the flame in this way 
supplied with the air which is 
necessary for it to burn, but the 
invisible gases which are given 
off are drawn over the small pieces 
of caustic soda, and the carbonic 
acid gas and water vapour ab- 
sorbed, whilst the remainder of 
the air (nitrogen) passes into the 
upper part of the bottle. On 
making the experiment, and again 
weighing the tubes and their con- 
tents, it will be seen that an m- 
crease in weight has occurred, and 
that nothing has been lost or de- 
stroyed. We can now understand 
that the burning of a candle is an example of chemical action. 
In the same way the burning of coal is due to the same chemical 
combination of the carbon and hydrogen of the coal with the 
oxygen of the air. 




Fig. II. 



What we have learnt 



In the third Lesson we have seen that sodium has the power of decom- 
posing water, Hberating hydrogen, and forming caustic soda, which turns 
red litmus blue. Water is decomposed or electrolised into the two gases, 
oxygen (one vol.) and hydrogen (two vols.) by passing through it a current 
of electricity from a battery. 

We have also learnt the difference between mechanical mixture and 
chemical combination. We have likewise learnt that in all the changes 
which the chemist brings about, no loss of matter occurs. He can neither 
create matter nor destroy it. We have seen how to prove that carbonic 
acid gas and water are produced when a candle burns, and that no loss, 
but an actual increase, in weight occurs, owing to the combination of the 
oxygen of the air with the materials, carbon and hydrogen, of the candle. 



EXERCISES ON WATER 



Exercises on Lesson ill 

1. Describe the action of sodium on water. 

2. How can you obtain the two gases oxygen and hydrogen from water ? 

3. Describe fully what happens wiien a candle burns, and mention any 
experiments which prove your assertions. 

4. Explain the difference between a mechanical mixture of two elements 
and a chemical compound of the same. 

5. What is meant when we say matter is indestructible ? 

6. How would you demonstrate the presence of (i) an acid, (2) an 
alkali dissolved in water ? 

7. Describe an experiment to show that when a candle burns, an increase 
in weight occurs. 

8. What is meant by electrolysis? 

9. Why must sodium always be kept under mineral oil ? 

10. The ancients considered water to be an element ; how can you show 
that it is not one ? 



LESSON IV 

ELEMENTS AND COMPOUNDS — SYMBOLS AND FOR- 
MULA—DISTRIBUTION OF THE ELEMENTS 

Elements and Compounds. — Having obtained a general 
idea of what Chemistry means, we will go a little farther into 
detail, using the knowledge we have already gained, to help us 
to understand some important principles. By experimenting 
on all kinds of substances, the chernist has found that whether 
they belong to earth, air, or sea ; to the animal, vegetable, or 
mineral kingdoms, they can be divided into two great classes, 
namely, Elements and Coinpomids . 

An element is a substance out of which nothing different 
has been obtained, or which has not been decomposed into 
two or more distinct and different substances. 

It is quite possible that some of the substances which we 
now call elements may, by further experiments, be split up into 
simpler bodies ; and such cases have, before now, occurred. 
The alkalis, potash and soda, were classed as elements until 
1807, when Sir Humphry Davy decomposed them and obtained 
the metals, potassium and sodium. 

The elements may be subdivided into metallic elements and 
non-metallic elements. 

A compound is a body formed by the chemical combination 
of two or more elements, and out of which two or more differ- 
ent elements can be obtained. 



OCCURRENCE OF THE ELEMENTS 



Common Elements, Metals, and Non-metals.* — 

Non-Metals. Metals {Solids). 

( Oxygen Iron 

Gases \ Hydrogen Silver 

Nitrogen ' Gold 

I Chlorine Lead 

Liquid Bromine Tin 

Iodine Zinc 

Carbon Nickel 

Solids Sulphur Mercury (Liquid) 

Phosphorus Sodium 

Arsenic Copper 

Silicon Aluminium 

Occurrence and Distribution of the Elements. — Com- 
paratively few of the elements occur in the free or uncombined 
state in nature. Amongst these are oxygen and nitrogen, these 
exist free and mechanicallly mixed together in the atmosphere. 
Sulphur occurs in the free state in large deposits in the neigh- 
bourhood of volcanoes, especially in Sicily. Carbon occurs 
free, in its purest natural form, crystallised as the diamond, 
in Brazil, South Africa, India, and other places. , It also 
occurs naturally in many places as graphite (commonly called 
black-lead although it contains no lead). Antimony, arsenic, 
copper, gold, silver, platinum, and mercury also occur in the 
free state. Iron is also found as large metallic masses in meteoric 
stones which have fallen to the earth from a source outside our 
atmosphere, these are seen and usually known as ^'shooting 
stars."" Although many of the elements occur in the free state 
in nature, they are more often found in the state of combination 
with other elements, e.g. the metals are usually found in metallic 
ores combined with one or more of the following elements — sul- 
phur, carbon, oxygen, silicon, chlorine, fluorine, phosphorus, etc. 

Hydrogen and oxygen occur combined as water, in addition 
to being essential constituents of all animal and vegetable bodies. 
Most organic compounds contain them, often combined also 
with nitrogen and sulphur. Man has only been able to pen- 
etrate a very short distance into the earth's solid crust, even 
in his deepest mines, so we cannot be certain about the com- 
position of the central portion of our globe. 

* A more complete list of the elements is given on p. 32. 



i.KSSON IV COMPOUND SUBSTANCES 23 

The various elements occur scattered very irregularly through- 
out the earth. Some are very abundant and occur widely 
distributed, whilst others have been found in such minute quan- 
tities, and in such rare fragments, that their properties have not 
yet been fully studied, e.g. oxygen occurs throughout the earth, 
sea, and air, in such quantities as to make up nearly half the 
weight of the earth's crust. 

The following figures show the approximate percentage com- 
position of the granitic rocks, of which the mass of the earth's 
crust is made up : — 



Oxygen 


44.0 to 48.7 


Calcium 


6.6 to 0.9 


Silicon 


22.8 to 36.2 


Magnesium . 


2.7 to o.i 


Aluminium . 


9.9 to 6.1 


Sodium 


2.4 to 2.5 


Iron . 


9.6 to 2.4 


Potassium . 


1.7 to 3.1 



Compounds. — The following are a few examples of com- 
pounds : carbonic acid gas, water, alcohol, sugar, caustic soda, 
oxide of phosphorus, phosphoric acid, oxide of mercury, lime, 
and chalk. 

The compounds formed by the combination of two or more 
elements are very often markedly different in their appearance 
and properties from the elements of which they are composed. 
We have already had some examples of this, and we will now 
take a few more from the compounds of the metals with non- 
metals. 

ExPT. 17. Copper in Bluestone. — Here, for instance, I 
have some bluestone or blue vitriol, wdiich contains copper, 
sulphur, and oxygen, although we find no indication of the fact 
on simply looking at it. 

It is easy, however, to prove that it contains copper. One 
has simply to dip one half of a clean steel knife blade in a 
sohitioii of the blue crystals in water, the immersed portion will 
at once be coated with a bright deposit of metallic copper. 

Mercury in Corrosive Sublimate. — Similarly this white 
crystalline substance called corrosive sublimate, a very poisonous 
body, contains the metal mercury. If we dissolve a little of it in 
water and dip a bright strip of copper, or a bright halfpenny, in 
the solution, the red metal becomes coated wdth a grey deposit 
of metallic mercury, which, if we rub it wdth the finger, will show 
a shining silvery layer of quicksilver. 



24 



DERIVATION OF SYMBOLS 



ExPT. 1 8. Silver and Lead contained in white Crys- 
talline Salts. — Again, this white crystalline solid substance, 
called lunar-caustic, contains silver. If some of it be dis- 
solved in a little distilled water in a test-tube, and a little grape 
sugar and caustic soda be added and the mixture warmed, we 
get the inside of the test-tube silvered by a thin coating of beau- 
tifully bright metallic silver. And again this other white com- 
pound (sugar of lead). If we dissolve a small quantity of it in 
water and hang a small piece of zinc in the solution in a tum- 
bler, it will yield, after a few hours, a beautiful tree-like growth 
of metallic lead called the lead tree. In these experiments a 
chemical reaction has taken place ; the iron has taken the 
place of the copper in the copper compound, and formed a 
new compound of iron, whilst metallic copper is liberated or 
set free in the pure state. Similarly, copper has replaced the 
mercury, and zinc the lead, new compounds of copper and zinc 
being formed, whilst the mercury and lead are set free in the 
metallic state. 

Symbols. — It is not always convenient to use the full names 
of the elements when we write them down, so chemists have 
devised Symbols which are to stand for them. These symbols 
are generally the first letter of the name or sometimes the first 
two letters ; and often where the first letter has once been used, 
a second element, whose name begins with the same letter, is 
denoted by the first letter or the first two letters of the Latin or 
Greek names. 

The following will serve as examples : — 



• Names. 


Symbols. 


Names. 


Symbols. 


Carbon 


C 


Sulphur 


S 


Calcium 


Ca 


Sodium 


Na (Natrium) 


Cobalt 


Co 


Silicon 


Si 


Copper 


Cu (Cuprum) 


Silver 


Kg (Argentum) 


Iodine 


I 


Nitrogen 


N 


Iron 


Fe (Ferrum) 


Nickel 


Ni 


Phosphorus 


P 


Magnesium 


Mg 


Potassium 


K (Kalium) 


Mercury 


Hg (Hvdrargvrum) 


Platinum 


Pt 


Hydrogen 


H 


Lead 


Pb (Plumbum) 


Oxygen 


O 



These symbols mean a very great deal more than the mere 
names of the elements as we shall see as we get further on. 



LESSON IV FORMULA AND EQUATIONS 25 



They are also used to denote compounds by placing the symbols 
close together. When two symbols are placed close together it 
means that the two elements denoted by them are cheinically 
co)iibi)icd and form a compound, e.g. : — 

HgO means Oxide of Mercury 

CuO " " Copper 

CaO " " Calcium (or Lime) 

MgO " " Magnesium 

and so on. 

Formulae and Equations. — These double symbols are 
called formulae (singular, formula), and it is by means of these 
formulae that the chemist expresses all his chemical reactions 
in eqiiatiojis. The following will serve as examples of equa- 
tions : — 

(Combination) Hg + O = HgO, 

meaning that mercury and oxygen have combined, as in Lavoi- 
sier's experiment, to form the red powder, oxide of mercury. 

(Combination) Cu + O = CuO 

denotes the formation of oxide of copper from metallic copper, 
and oxygen, e.g. when Cu is very strongly heated in air, it be- 
comes covered with black scales of oxide of copper. 

(Decomposition) HgO = Hg + 0. 

This denotes that the red powder has been again decomposed 
into its elements. It will be noticed that on the left-hand side 
is placed the condition of the elements or compounds before the 
chemical reaction (whether combination or decomposition) has 
taken place, whilst on the right-hand side are placed the products 
or ultimate result of the chemical action, and this order is invari- 
ably followed, thus : — 

(Combination) Hg -f O = HgO. 
(Decomposition) HgO = Hg + O. 

Matter is indestructible, it can neither be created nor destroyed. 
It can only change its form by entering into new combinations ; 



26 SUMMARY AND EXERCISES PART I 

therefore, there is always the sai)ie weight of materials on each 
side of an equation^ e.g. in the above equations, the weight of 
mercury and the weight of oxygen added together, are exactly 
equal to the weight of the oxide of mercury. 

What we have learnt 

In our fourth Lesson we have learnt that all substances maybe classified 
as elements and compounds. Elements are substances which have not been 
chemically decomposed into two or more different bodies. Compounds 
are substances formed by the union or combination of two or more ele- 
ments. We have seen that compounds differ in the most marked manner 
from the elements of which they are composed. The elements are divided 
into metals and non-metals. Symbols are used to denote the names of the 
elements, whilst formulae are used to denote the composition of compounds. 
Equations are used to exhibit chemical reactions by means of symbols 
and formulae. 



Exercises on Lesson IV 

1. Give five examples each, of elements and compounds. 

2. Give four gaseous and four solid non-metals. 

3. What is a compound ? 

4. Write out three formulae and three equations. 

5. How can you show that blue vitriol contains copper, and that corro- 
sive sublimate contains mercury ? 

6. What is the difference in meaning between the ecj[uation HgO = Hg 
+ O and that Hg + O - HgO ? 

7. What elements are denoted by the following symbols, Fe, I, Ni, Mg, 
Hg, Ag, Na, S, Si, K, Pt, Pb, Co, Cu ? 

8. What are the symbols for mercury, lead, sulphur, silver, silicon, 
sodium, iron, iodine ? 



LESSON V 

COMBIXATIOX IX DEFINITE AND AIULTIPLE PROPOR- 
TIONS — DALTON-S ATOMIC THEORY — ATOMIC .\ND 
MOLECULAR WEIGHTS — CALCULATIONS 

Constancy of Chemical Composition. — These symbols, 
formulae, and equations have a far ^vicler meaning than is in- 
dicated in the preceding lesson : and to understand what that 
meaning is, we must try to master the laws of combination in 
definite and multiple proportions. 

To understand this most important subject let us begin with 
a simple illustration. By long experience and many experi- 
ments, chemists have found that every chemical compound has 
a fixed and definite composition. No matter how it is made or 
where obtained, any given chemical compound, whether found 
in nature or artificially made, if carefully analysed is always 
found to contain exactly the same proportion by weight of the 
elements of which it is composed. Thus for example, if we take 
215 parts of the red powder, oxide of mercury, and then decom- 
pose it by heat, and are careful not to alloAv any of the mer- 
cury to escape as vapour, we obtain 199 parts of liquid metallic 
mercury, whilst 16 parts by weight of gaseous oxygen escape. 
Similarly, if we take 24 parts by weight of the metal mag- 
nesium and burn it carefully so as not to lose any of the white 
fumes, we get 40 parts by weight of oxide of magnesium, and 
we must remember that in burning, magnesium is combining 
with oxygen. Again, if we take 79 parts by weight of black 
oxide of copper and abstract all its oxygen from it (as we can 
do bv heatins: it strono^lv in a current of hvdrosren eras, the 



28 RELATIONS BETWEEN THE part I 

hydrogen combining with the oxygen to form water which es- 
capes as steam, leaving metallic copper behind) we find that 
63 parts of metallic copper remain, and 16 parts of oxygen 
have disappeared. Therefore, these oxides have the following 
composition by weight : — 

Mercury 199 Magnesium 24 Copper 63 

Oxygen 16 Oxygen 16 Oxygen 16 

Oxide of Mercury 215 Oxide of Magnesium 40 Oxide of Copper 79 

Relations bet"ween the Combining Weights. — If now we 

compare the sidphides of these metals (the sulphides are com- 
pounds formed by the combination of the metals with sulphur, 
just as oxides are the compounds formed with oxygen) we find 
the following are the proportions : — 

Mercury 199 Magnesium 24 Copper 63 

Sulphur 32 Sulphur 32 Sulphur 32 

Sulphide of Mercury 231 Sulphide of Magnesium 56 Sulphide of Copper 95 

Similarly with the chlorides — combinations of the metals with 
chlorine. We have the following numbers : — 

Mercury 199 Magnesium 24 Copper 63 

Chlorine 70 Chlorine 70 Chlorine 70 

Chloride of Mercury 269 Chloride of iMagnesium 94 Chloride of Copper 133 

Here we see a remarkable relation existing between the given 
weights of the three metals and a constant weight of another 
element, this constant being a different one for each element 
taken. Thus if we always take the same proportions of the 
metals, namely, 199 parts of mercury, 24 parts of magnesium, 
and 63 parts of copper, we find that a constant weight of some 
other element, 16 of oxygen, 32 of sulphur, or 70 of chlorine, 
will be required to combine with the metals to form a definite 
compound. 

This remarkable relation was, however, not at once seen by 
chemists ; they were accustomed to express their cliemical 
compositions by giving the number of parts of each element 
required to make up 100, as in the figures below, and in the 



LESSON V COMBINING WEIGHTS 2^ 

numbers expressing percentage composition, no definite relation 
seemed to exist. 

Up to the time of John Dalton (1808), there was no satis- 
factory manner of accounting for this remarkable constancy of 
the exact proportions in which the elements were found com- 
bined together in compounds. Apparent exceptions to the 
uniform constancy of composition, however, led John Dalton to 
formulate a theory which accounted in a completely satisfactory 
manner for all the ascertained facts, the importance of which 
theory cannot be over-estimated, for it forms one of the prin- 
cipal foundation stones of the exact science of chemistry. 
Dalton was acquainted with several cases in which two ele- 
ments combined together in two different proportions, e.g. 
carbon and oxygen, and carbon and hydrogen each formed two 
compounds thus : — 

(I) (2) (I) (2) 

Carbon 42.86 27.27 Carbon 85.68 74.95 

Oxygen 57.14 72.73 Hydrogen 14.32 25.05 

100.00 100.00 100.00 100.00 



These figures show no definite relation one to the other, but 
Dalton asked himself what was the relation of one element, say 
oxygen, in both compounds when the other element (carbon) 
remains constant? In other words, how much oxygen in each 
case combines with the unit weight of carbon. Let us calculate 
this. Divide 42.86 and 57.14 by 42.86; we get i of carbon and 
1.33 of oxygen. Then divide 27.27 and 72.73 by 27.27 ; we get 
I of carbon and 2.66 of oxygen or : — 

1st Compound. 2nd Compound. 

Carbon = i Carbon = i 

Oxygen = 1.33 Oxygen = 2.66 

Now we see a definite relation between the two quantities 
of oxygen, for we find that for the same weight of carbon, the 
oxygen in the second compound is exactly donble that in the 
first. In the same way the hydrogen compounds contain 

1st Compound. 2nd Compound. 

Carbon = i Carbon = i 

Hydrogen = 0.167 Hydrogen = 0.334; 



30 THE FIVE OXIDES OF NITROGEN part i 

or in the second compound, for the same weight of carbon, the 
amount of hydrogen is exactly double the amount in the first 
compound. 

Again a series of five compounds of nitrogen and oxygen 
are now known (which we shall study more fully afterwards). 
Looking at their percentage chemical composition, no particular 
relation or ratio between the figures is noticed. 

(I) (2) (3) (4) (5) 

Nitrogen 63.6 46.6 36.8 30.4 25.9 

Oxygen 36.4 53.4 63.2 69.6 74.1 



If, however, we calculate the weight of oxygen combined with 
a constant weight of nitrogen (say 27.8 parts of nitrogen,*) we 
find, as in the foregoing examples, that oxygen is present in the 
simple ratios 1:2:3:4:5, thus : — 





(1) 


(2) 


(3) 


(4) 


(5) 


Nitrogen 


27.8 


27.8 


27.8 


27.8 


27.8* 


Oxygen 


15-9 


31.8 


47-7 


63.6 


79-5 



There are a very great many examples of two elements 
forming several compounds in this manner ; but taking a con- 
stant weight of one element, the other is always present in 
weights which are either simple multiples of the lowest weight, or 
which have a very simple relation one to the other, such as 

1:2 2:3 3:5 

These facts are expressed in Dalton's Laws of co7nbmation in 
Definite and Multiple proportions. 

Dalton's Atomic Theory. — To explain ihese facts, Dalton 
proposed the following theory known as Dalton's Atomic 
Theory, which is of such importance that the student must 
understand it thoroughly. He assumed, following the views 
of the old Greek philosophers, that all matter is made up 
of exceedingly small invisible particles called atoms (from 
a privative, and re/xi/w I cut), that these atoms cannot be 
divided into two smaller portions, in other words that the 

* 27.8 is twice the atomic weight of nitrogen (see p. 33). 



LESSON V DALTON'S ATOMIC THEORY 31 

atom is the smallest particle of matter which can exist. He 
further assumed that chemical combination consists in the 
joining together of the atoms of the combining elements 
to form a group of atoms or a inolenile * which are bound 
to each other by their mutual chemical attraction. Further, 
that all the atoms of the same element are of exactly the 
same weight, and possess the same properties, and that atoms 
of different elements are different in weight, and possess other 
properties. 

We can now see w^hy chemical compounds possess a constant 
composition, because each molecule of a compound is exactly 
like every other molecule, it is composed of the same number 
of atoms of the same kinds and possessing the same weights. 
Hence the whole mass of a substance consisting of millions of 
molecules will contain just the same proportion of the com- 
ponent elements as a single molecule does. Further, suppose 
we have a molecule, say of one atom of carbon and one atom 
of oxygen. To make a compound containing more oxygen, we 
must add at least one more atom, we cannot add part of an 
atom, but the addition of another atom of oxygen to every 
molecule exactly doubles the proportion of oxygen, the carbon 
remaining constant, and this is in acc.ordance with the facts. 
Similarly with the five compounds of nitrogen and oxygen. 
The first compound only contains one atom of oxygen in the 
molecule, the second two, the third three, and so on ; the 
fifth compound contains five atoms of oxygen in the molecule. 
We cannot have any intermediate compounds according to 
Dalton's atomic theory, and this harmonises with the facts, for 
it has hitherto been found impossible to prepare any intermediate 
compounds of nitrogen and oxygen. 

Evidently then it is possible to find out by chemical analysis, 
not only the percentage composition of a body, but also the 
relative weights of the atoms. This has been done for all the 
elements at present known, and as hydrogen is the lightest 
substance known, the ato7?t of hydrogen is taken as the tinit of 
weight. Therefore we say that the atomic wxight or the weight 

* Sir William Thomson (now Lord Kelvin) has calculated that if a 
drop of water could be magnified so as to appear as large as the earth, the 
molecules would appear about as large as cricket balls. How many 
millions of millions of molecules then must there be in a drop of water ? 



32 ATOMIC WEIGHTS 



of the atom of hydrogen is i or H = i. We have seen that 
sodium displaces hydrogen from water ; 

H.p + Na = NaOH + H. 

If we find out how many grams of sodium are required to 
displace one gram of hydrogen, we obtain the atomic w^eight of 
sodium. This is found to be 22.9. Caustic soda, the first 
product of the action of sodium on water, is found still to contain 
as much hydrogen as the water has lost. Therefore the molecule 
of water must contain two atoms of hydrogen. This is con- 
firmed by the fact that in the electrolysis of water we get twice 
as large a volume of hydrogen as of oxygen. Therefore we 
write the formula of water H2O, the little 2 is placed beneath and 
to the right of the symbol to denote two atoms. 

Water is found to be composed of 2 parts by weight of 
hydrogen and 15.9 parts by weight of oxygen, so we say the 
atomic weight of oxygen is 15.9, and oxygen gas is found to 
be exactly 15.9 times as heavy as hydrogen gas. 

By similar methods the atoDiic weights of all the known 
elements have been found. But all the minute precautions 
and all the abstruse considerations which are taken into account 
in finding them cannot be entered into here. Suffice it to say 
that as the atomic weights of the elements are of the utmost 
importance and will be constantly needed in subsequent lessons, 
it will be well here to give a list of them. Only those elements 
which are most important, or which are most plentiful and 
useful to man are given. 

The names printed in large capital italics are the non-metals ; 
those in small Roman capitals are the metals. 

List of Elements and their Atomic Weights. — 



Names. 


Symbols. 


Atomic Weights.* 


Aluminium 


Al 


26.9 


Antimony 


Sb (Stibium) 


119.4 


ARSENIC 


As 


74.4 


Barium 


Ba 


136.0 


Bismuth 


Bi 


206.4 



* The above atomic weights are given to one place of decimals only, 
as being sufficiently accurate for use in elementary work. 



LESSON V 



ATOMIC WEIGHTS 



33 



Names. 


Symbols. 


Atomic Weights. 


BORON 


B 


10.7 


BROMINE 


Br 


794 


Cadmium 


Cd 


111.3 


Calcium 


Ca 


39.7 


CARBON 


C 


11.9 


CHLORINE 


Cl 


35-2 


Chromium 


Cr 


51-9 


Cobalt 


Co 


58.6 


Copper 


Cu 


62.8 


FL UORINE 


F 


18.9 


Gold 


Au (Aurum) 


195.7 


HYDROGEN 


H 


I.O 


IODINE 


I 


125.9 


Iron 


Fe (Ferrum) 


55.6 


Lead 


Pb (Plumbum) 


2054 


Magnesium 


Ug 


24.2 


Manganese 


Mn 


54-6 


Mercury 


Hg (Hydrargyrum) 


198.9 


Nickel 


Ni 


58.6 


NITROGEN 


N 


13-9 


OXYGEN 


O 


15-9 


PHOSPHORUS 


P 


30.8 


Platinum 


Pt 


193.3 


Potassium 


K (Kalium) 


38.9 


Silver 


Ag (Argentum) 


107. 1 


SILICON 


Si 


28.2 


Sodium 


Na (Natrium) 


22.9 


SULPHUR 


S 


31.8 


Tin 


Sn (Stannum) 


1 17. 2 


Zinc 


Zn 


65.0 



There are other elements (making in all about 70), which 
are rare or not of such importance as those named. 

We are now in a position to understand how necessary it 
is to know the atomic weights of the elements, and also the 
exact composition of any chemical compound. We now see 
that a chemical formula not only expresses the nature of the 
elements composing the compound, but also gives its quantitative 
composition or the proportions by weight in which each element 
is present. In many cases chemical compounds contain two, 
three, or more atoms of the same element, and when this 
is the case we must modify the simple formulae which we have 
previously used by adding a small figure just beneath and to 
the right of the symbol of such element. 

The symbol of an element represents one atojn and the formula 
of a compound represents o)ie molecule. 



34 CALCULATION OF QUANTITIES part i 

In the five compounds of nitrogen and oxygen, for example, 
we see on looking at the table of atomic weights that N = 13.9 
and O = 15.9; but the first compound contains N 27.8 and O 
15.9, therefore there are two atoms of nitrogen and one atom of 
oxygen in the molecule, and the formulae of the five compounds 
may be expressed as N^O, N^O^,* N^^Og, NgO^,* and NgO^, 

The carbon and oxygen compounds are CO and CO2 
The carbon and hydrogen '* CH^f and CH^ 

Calculations. — Let us now take a few examples of the 
way in which these atomic weights are used in calculations of 
weights of materials used in chemical reactions. Potassium 
chlorate is a white solid body having the formula KClOo. This, 
when strongly heated, decomposes, and gives off all its oxygen, 
leaving behind a compound called potassium chloride (KCl). 

Example. — Find out how many ounces of potassium chloride 
remain, on completely decomposing 100 ounces of KClOo, 
and also calculate the percentage composition of the latter 
compound. 

Directions. — i. If a chemical action occurs, always write out the 
equation expressing it. 

2. Write down the symbols of the various elements of the compounds 
concerned in the calculation under each other, and place the atomic 
weights on their right, not forgetting to multiply the atomic weights by the 
number of atoms where more than one atom of any element is present as 
in KCIO3, Oxygen = 15.9 X 3. Next add up the numbers to obtain the 
molecular weight, or weight of a single molecule. In the third term of the 
proportion sum that follows place the given weight of substance. In the first 
term place the atomic or molecular weight of the substance occurring in 
the third term. In the second term place the atomic or molecular weight 
of the substance sought. 

KCIO3 = KCl + 3O. 

Os = 47'7 47-7 = 30' 

121. 8 121. S = KClOq. 



* We shall see later that these formulae should be NO and NOo, but 
the ratio between the nitrogen and oxygen is not thereby altered, so the 
argument is unaffected. 

t Similarly, the formula of this compound is really C.2H4 (see p. 232),' 
but the ratio between the carbon and hydrogen is not altered. 



LESSON V IN CHEMICAL REACTIONS 35 



Therefore 121.8 parts (grams, ounces, lbs., or tons) of KCIO3 yield 47.7 
parts of oxygen, and 74.1 parts of KCl. Tlien, by simple proportion as 
above described, we have 

121. 8 : 74.1 : : 100 : Ans. 
74.1 X 100 _ 7410 



121. 8 121. 8 



: 60.84. Ans. 



Answer 60.84 parts of potassium chloride remain. 



Second Method. —It may perhaps be easier for the younger readers 
to work out the above calculations by the method of reduction to unity^ 
instead of by proportion as explained above. Let us take the first calcula- 
tion on p. 34. In the second method we simply ask — 

If 121.8 parts of potassium chlorate yield 74.1 parts of potassium 
chloride, what will i part yield ? Evidently we must divide 74.1 by 121.8^ 

or^ii^. Then, if i part yields -Ziil, what will 100 yield? Evidently 
121.8 121. 8 

74-1 74-1 X io<^ 7410 u- u • .^ u . X I, 

^ X 100, ov-^-^ — , or -, which is exactly what we come to by 

121.8 121.8 121.8 ^ ^ 

the first method. Or 

If 121. 8 gives 74.1 

„ ^ 

121.8 

74-1 
/. 100 -^-^ X 100 



All the other calculations can be worked oiit in the same way. 

If we calculate in the same way how many ounces of oxygen the com- 
pound has lost, the two answers should add up to 100. Let us see if 
they do. 

121. 8 : 47.7 : : 100 : Ans. 

47-7 X 100 4770 ^^. .„, 

= o — = n = 39-IO- Ans. 

121. 8 121. 8 ^ 

Answer 39.16 parts of oxygen given off. 

.*. Oxygen given off 39.16% 

Potassium chloride left behind 60.84% 

Potassium chlorate decomposed . . . 100.00 



If we find the percentage of K, and subtract it from 60.84 ^o obtain the 
chlorine, we obtain the percentage composition of KCIO3. Thus 

121. 8 : 38.9 : : 100 : Ans. 

38.9 X 100 3890 . 

^ — = ^^-^ = 31.94% of potassium. 
121.8 121.8 ^ ^^'^ ^ 



36 PERCENTAGE COMPOSITION part I 



The percentage composition is therefore 



K . . . 

CI ... . 28.90% 

0,s . . . . 39-i6 % 

KCIOq . . . 100.00 



3I.94%|kc1 = 6o".84%. 
28.90% ) 



Further Examples.*— (i) What weight of liquid mercury can be 
obtained by the decomposition of 4 ounces of oxide of mercury, HgO ? 

HgO = Hg + O. 

Hg 199.0 214.9 • 199 • • 4 * Ans. 

^ _1S:9 ^ ^99 X 4 ^ Ans. 

HgO 214.9 214.9 

(2) From the following equation calculate what weight of sodium will be 
required completely to decompose 2 lbs. of water. 



H2(i X 2) 
O 



Na H- H2O = NaOH + H. 
~ isig I ^-^^ ^ ^7-9. Na =22.9. 



22.9 X 2 , 
17.9 : 22.9 : : 2 = — z ^ ^ns. 

17.9 

(3) How much phosphoric acid, H3PO4, will be produced by burning 2 
ounces of phosphorus in a closed bell jar over water ? 

P-2 + O5 = P.2O5. P.2O5 + 3H2O = 2H3PO4. 

Two atoms of phosphorus yield 2 molecules of phosphoric acid, there- 
fore I atom of P yields i molecule of H3PO4. 

HgCi X 3) - 3.0 

P = 30.8 30.8 : 97.4 : : 2 

04(15.9 X 4) =63.6 

H3PO4 =^974 ^974x2^^^^^ 
30.8 

(4) How many ounces of water will be required to form phosphoric acid, 
H3PO4, with 6 ounces of PoOg ? 

P.205 + 3H.p = 2H3P04. 



*The student should work out these calculations himself. 



LESSON V SUMMARY AND EXERCISES 37 

Therefore 3 molecules of water are required for each molecule of oxide 
of phosphorus, and we must not omit to multiply by 3. 

3H2O = (2 + 15.9) X 3 = 53.7 
P2 (30.8 X 2) = 61.6 
O5 (-5-9x5) - 79-5 141.1 -SS'? • :6 

P,05 -141.1 = 5^:Z_^ = Ans. 

=^ 141. 1 

Calculation of Formulae. — As the converse of the first example 
given on p. 34, if we know the percentage composition of a compound it 
is easy to calculate its simplest formula. For if we divide the percentage 
of each element by its atomic weight and divide each number so obtained 
by the lowest, we obtain the ratio between the number of the atoms, e.£: — 

K = 58.9 •3i.94%K ^ = 0.821 l"^ = i = K. 

CI = 35.2 28.90% CI ??^ = 0.821 'l^^i^CI. 
^^ 35.2 .821 

0=^15.9 39.16% o ^ = 2.463 ^^3-03. 

Therefore the formula of potassium chlorate is KCIO3. 
What we have learnt 

In the fifth Lesson we have learnt a wider meaning of symbols, formulae, 
and equations. Constancy of composition of all compounds. Combining 
weights and the relations between them. Combination in multiple pro- 
portions, illustrated by oxides and hydrides of carbon, and the five oxides 
of nitrogen. Explanation of above facts by Dalton's Atomic Theory. 
Meaning of atomic weights, of which a list is given. Method of using 
atomic weights in the calculation of percentage compositions, and the 
weights of materials taking part in chemical reactions. 

Exercises on Lesson V 

1. What is meant by the constancy of composition of chemical com- 
pounds ? 

2. A chemical compound is found to consist of hydrogen 3.08%, phos- 
phorus 31.62%, oxygen 65.30%. What is its simplest formula ? 

3. Give some examples of chemical combination in multiple proportions ; 
why were not the relations between the combining weights at once seen ? 

4. Give an account of Dalton's Atomic Theory. 

5. Work out the numerical examples given in the lesson. 

6. Calculate the percentage composition of manganese dioxide, MnO^, 
potassium nitrate, KNO3, and sodium sulphate, Na2S04. 

7. How much sodium shall I require in order to obtain 400 grams of 
pure caustic soda ? 

8. From the percentage compositions of the oxides of nitrogen given on 
p. 30, calculate their simplest formulae. 



LESSON VI 
PHYSICAL MEASUREMENTS 

STANDARDS OF LENGTH, VOLUME, AND WEIGHT USED 
IN CHEMICAL EXPERIMENTS — METHODS OF WEIGH- 
ING AND MEASURING -THE METRIC SYSTEM— THE 
THERMOMETER— MEASUREMENTS OF TEMPERATURES 
— CONVERSION OF THERMOMETRIC SCALES 

Standards of Length. — The standard unit of length in 
the metric system, which is now generally used for scien- 
tific purposes, is the metre (m.). This is divided into tenths 
(decimetres), and into hundredths (centimetres) and thou- 

123456789 10 

IiimJiiiiIiimIimiIiiiiIiiiiIiimIiiiiIiiiiIiiiiIiiiiIimiIiiiiIiiiiiiiiH^ 

Fig. 12. 

sandth (millimetres). Fig. 12 is an exact copy of a deci- 
metre scale, divided into 10 centimetres and into 100 milli- 
metres. The scale is, therefore, one-tenth of a metre in length 
(decimetre). 



Thus we have 

I decimetre (dcm.) = o.i m. 
I centimetre (cm.) =0.01 " 
I millimetre (mm.) = o.ooi " 



10 dcm. = I m. 

100 cm. = I " 

1000 mm. = I " 



10 m. = I deca-m. 
100 m. = I hecto-m. 
1000 m. = I kilo-rp. 

For the relation between the English and metric systems, see Appendix. 

The Measures of Area are easily obtained from the 

38 



LESSON VI THE METRIC SYSTEM 39 

standards of length. Thus square metre, square decimetre, 
square centimetre, square millimetre. Fig. 13 is exactly a 
square centimetre, a cube on this base would be a cubic centi- 
metre, the unit of volume, which, as we shall see pre- 
sently, is closely related to the gram, the unit of weight. 
The Standards of Volume. — The unit of volume 
is the cubic centimetre, but one cubic decimetre = 1000 



cubic centimetres is the common unit of volume, and ^^' ^^' 
IS called a //trey and we have decalitre = 10 litres, hectalitre, 100 
litres, kilolitre, 1000 litres. 



I cubic centimetre = unit 


10 ]itres = deca litre 


100 cc. —decilitre 


100 " =hecta " 


1000 cc. == litre 


1000 " =kilo " 



The Standards of Weight. — The standard of weight is 
the weight of a cubic centimetre of distilled water at its point of 
maximum density, 4" C.,* weighed at Paris,t and this weight is 
called a gramme, or (English) gram. The common unit is a 
kilogram, or 1000 grams. 



I gram= 10 decigrams 
= 100 centigrams 
= 1000 milligrams 



10 gr.= i deca-gr. 
100 " =1 hecto-gr. 
1000 " =^1 kilo-gr. 



The Chemical Balance. { — In all important chemical 

* The reason for taking 4° is explained on p. 102. 

t The true weight of a body varies at different places on the earth's 
surface, because the attraction of gravitation is not uniform. Owing to 
the earth being flattened at the poles and bulged out at the equator, and 
also owii-'.g to its rotation, a substance (say a gram weight) weighed on a 
spring balance at the poles will be nearer the centre of the earth (at which 
point the attractive force of gravitation may be supposed to act) and will 
weigh more than if weighed at the equator. Hence, to be strictly ac- 
curate in such an important matter as defining the standard of weight, it is 
necessary to state the place at which the standard is correct, e.^". A gram 
weight will weigh exactly a gram at Paris, but more than a gram at London, 
and still more at the poles. But although its weight varies, its mass, or the 
quantity of matter it contains, will be the same everywhere — and if it is 
desired to eliminate the varying force of gravity, we speak of the mass of a 
body instead of its weight. Its mass is constant, but its weight depends 
upon where it is weighed, because the attractive force which gives bodies 
weight, viz. gravitation, varies at different places. 

\ For further information about the balance and methods of weighing 
consult Stewart and Gee's Practical Physics, vol. i., from which book Figs. 
14-16, 19-21, and 59 have been taken. 



40 



THE CHEMICAL BALANCE 



experiments which have for their object' the determination of 
the quantitative chemical composition of substances, the chemist 
makes use of delicate instruments for measuring and weighing 
all the materials which he uses. Perhaps the most important 
is the chemical balance, which consists essentially of (i) the 
beam b^ or arms of the balance, which may be of brass or steel 
made in some such pattern as is shown in Fig. 15, so as to 
be as light and yet as rigid as possible. At each end are 
two knife edges of agate^ k^ turned upwards, and in the centre 




Fig. 14. 

another knife edge, K, turned downw^ards. The centre knife 
edge rests on a perfectly smooth plate of agate fixed on the 
pillar P of the balance. At the centre of the beam, and at 
right angles to it, pointing downwards, is a long pointer, p, which 
moves over a graduated scale. When the two sides are exactly 
balanced, the pointer rests at zero, the centre of the scale, 
whilst if either side be heavier than the other the pointer will 
move to the opposite side. There is a small vane v which 
may be moved to one side or the other, so as to adjust the 



LESSON VI 



THE CHEMICAL BALANCE 



41 



weight on each side very accurately, in order that when the 
pans are unloaded the pointer may vibrate equally on each 
side of the zero point. The two pans are suspended on the 
agate knife edges, which work with a minimum of friction on 




Fig. IS- 

agate planes. When not in use the pans may be lifted off the 
beam, and the beam from the pillar so as to support them on 
an independent arm, by a single turn of the large screw M. 
This preserves the knife edges from wearing, due to any vibration 
taking place when the balance is not in actual use. 



Use of the Rider in exact weig'hing'. — The beam is divided on 
each side into tenths and twentieths of its length measured between the 
two knife edges and a very small weight called a rider 
(Fig. 16), because it is placed striding across the beam, 
may be moved along by means of the rod d, to any of 
the positions marked on the beam. This rider generally 
weighs a centigram, but if placed at ^^ the arm-length 
from the centre its effect is just ^-^ of what it would be 
if placed on the pan, just as in the lever, Fig, 17, which 
would be in equilibrium with a i lb. weight at one 
end, and a 10 lb. weight at the other if the distances 
between the knife edges were as 10 to i. The 10 lb. 
weight has the same effect at unit distance as a i lb. weight has at 10 
times the distance from the central support. Evidently then if we place 
our rider at a position on the beam marked j.^^, and the balance is in 
equilibrium, it denotes a weight on the opposite pan of 7.3 tenths of a 
centigram, or .0073 gram. It would not be convenient to make weights 
so small as .0001 gram, or the ten thousandth of a gram, yet by means of 



A 



Fig. 16. 




42 THE CHEMICAL BALANCE part I 

the rider it is easy to detect accurately on a good balance a difference 
equal to this slight amount by means of a weight loo times as heavy (and 
yet this is only a small piece of twisted aluminium wire). 

When each pan contains a weight equal to loo grams, a 
good balance will indicate a difference of a tenth of a milligram, 
or .0001 gr., therefore it will indicate a diiference of a toooooo? 

a millionth part of 
.,g J the load. For mak- 
ing delicate weigh- 
ings the balance 
must be enclosed 
Fig. 17. in a glass case, so 

as to protect it frorn 
air currents, which would have the effect of disturbing the 
equilibrium. To protect the balance from the effects of mois- 
ture and rust, especially if it has steel arms, it is usual to place 
inside the balance case, a funnel containing dry calcium chloride, 
which rests in an empty bottle and which wdll absorb the 
moisture readily, and preserve a dry interior. 

Conditions of Sensitiveness and Equilibrium. — The 
conditions which must be observed to make the balance sensitive 
are (i) the beam to be as light and as long as possible 
consistent with rigidity, and (2) the centre of gravity of the 
whole (beam, pans, and weight) to be below, and as near as 
possible to the centre of suspension. 

The whole w^eight (of beam, pans, and weights) may be supposed to 
act as its centre of gravity E, the turning weight must raise this weight to 
^ at a distance Sa from the 
point of support, so as to j^^ 

be in equilibrium. It is iv ^^-- 

evident that if the centre 
of gravity of the whole be 
at E', twice as far away 
from the axis of support, it 
will require twice as great 

a w^eight to give the same '' • 

deflection of the balance, 

because the weight must ^^* ■^ * 

then be moved twice as 

far away (S<^) from the point of support. The gravity-bob g at the 
centre of the beam is made to turn up and down, so as to adjust 
the point of the centre of gravity as near as possible to the centre 



A 


^-^ S A' 


V 




E ^-^^ 



{ 



LESSON VI 



METHOD OF WEIGHING 



43 



of support. Then the whole weight acting at E must be moved an ex- 
ceedingly small distance Sa from the point of support. If the centre of 
gravity E comes adove the point S, the system is in a position of unstable 
equilibrmm, and the beam would overturn with the least touch; whilst 
if S and E are comcident, the beam would be in equilibrium in any position 
(neutral equilibrmm). The accuracy depends on the two arms being of 
exactly the same length, and this may be tested by weighing a substance 
first on one pan and then on the other — if both weighings agree, then the 
arms are equal. WxA^W xA'; therefore if W = W, then A = A , 
and since W and W are equal, they may be interchanged, then W X A = 
W X A'. 

If the two weighings do not agree, the arms are unequal, and the true 
weight is found by taking the square root of the product of the two apparent 
weights. 

Method of Weighing — the Weights. — The set of 

weights used for exact work in chemistry usually consists of 
500, 200, TOO, 100 grams; 50, 20, 10, 10 grams; 5,-2, i, i, 
I. grams, making up a 
kilogram; then .5, .2, 
.1, .1, .05, .02, .01, .01, 
and a couple of riders, 
each equal to a centi- 
gram ; sometimes .005, 
.002, .001, .001 are pre- 
sent in the set, but are 
seldom used, the rider 
doing all that is neces- 
sary below o.oi- The 
substance to be weighed 
is placed in the left pan, ^*^- ^9- 

and the nearest weight is placed on the right ; if this is too 
heavy, the next lower is taken, and so on until it is balanced 
to a centigram ; the final adjustment is made by means of the 
rider, which is moved along the arm until the pointer vibrates 
equally on both sides of the zero point. Suppose the weights 
are 50, 2, i, .2, .05, and the rider at 7.3, the weight of the 
substance is 53.2573 grams. 

How to measure Volumes of Liquids exactly. — 
For this purpose marked measuring flasks are used (Fig. 20) ; 
these have a circular mark e etched on the neck, and must be 
filled exactly to this mark, at a standard temperature. For 
withdrawing a definite volume of a liquid, pipettes are used 




44 



MEASUREMENTS OF VOLUME 



PART I 




(Fig. 20), the point of the pipette is immersed in the 
liquid, which is sucked up by the mouth as far as possible. 
The finger is then quickly placed on the top of the 
pipette, and carefully released just a little to allow air to 
enter ; this enables the liquid to flow out until the mark 
a on the stem is reached, 
w^Tcn the desired volume is 
obtained. When intermediate 
volumes are required, gradu- 
ated cylinders, pipettes, or 
burettes (Fig. 21) are em- 
ployed. 

Measurement of Temper- 
ature, the Thermometer. — 
As all bodies — solid, liquid, 
and gaseous — expand when 
heated, it is necessary to be 
able to ascertain their tem- z\ 
perature, so that their volume 
at a constant temperature may 
be known. 

ExPT. 19. — Here is a large 
glass bulb filled with a 
coloured liquid to a mark on 
the stem at the ordinary temperature of the air. On 
plunging this into a large beaker of hot water, W'C notice 
first an immediate but slight fall in the level of the ^^^" ^^' 
liquid, and directly afterwards the liquid begins to rise again, 
and, passing the first point, rises high in the stem. Fig. 22 
shows that the first effect of the heat was to expand the solid 
glass bulb and make it larger, and hence the fall, for to fill it 
to the same mark more liquid is required. The heat, how- 
ever, soon reaches the liquid, and it too begins to expand ; but 
it expands 77mch 7nore than the solid glass, hence it not only 
fills the enlarged bulb to the same place from which it started, 
but it rises higher and higher, to the second mark on the 
stem. 

ExPT. 20. — Similarly, that gases expand by heat may be 
shown by fitting up a large flask with a delivery tube dipping 
into a pneumatic trough. By applying a Bunsen burner to the 



Fig. 20. 



LESSON VI 



THE THERMOMETER 



45 



outside of the flask, the contained air expands so much that 
a quantity of it escapes and ascends into the gas cylinder. 
On taking away the lamp, how- 
ever, the air soon cools down 
again, and in cooling contracts^ 
and the water rises up the delivery 
tube into the neck of the flask to 
take the place of the air which has 
escaped (Fig. 23). 

Hence it is quite evident that 
if we want to know accurately the 
volume of any solid, liquid, or gas, 
we must know their temperature, 
for they occupy different volumes 
at different temperatures. 

The thermometer is used for 
measuring temperature. It con- 
sists of a thick glass tube with an 
extremely fine bore, terminating 
in a spherical or cylindrical bulb, 
containing a liquid, generally mer- 
cury or alcohol. Mercury is most 
often used, because it expands uni- 
formly, and because it solidifies at 
an extremely low temperature and boils at a very high tempera- 
ture, in other words, a mercurial thermometer has a very great 
range. The thermometer is made in the following manner. 
The open end is enlarged into a little cup, which is filled with 
mercury ; the bulb is heated by a Bunsen flame ; this expels a 
portion of the air, and on cooling, a portion of mercury enters 
the bulb (Fig. 24). By alternate heating and cooling, ^// the 
air is expelled, and the bulb and stem completely filled with mer- 
cury. It is then heated to the boiling point of mercury, and 
when it has cooled to the highest temperature which it is desired 
to measure, the end is sealed up by fusing the glass in a blow- 
pipe. When the mercury cools it recedes in the tube and no air 
can enter. 

Graduation of the Thermometer. — The next thing is to 
gradiiate the thermometer, and for this purpose two points 
are taken as starting points from which to graduate the whole 




Fig. 22. 



46 



GRADUATION OF THERMOMETERS 



stem. The first is obtained by plunging the thermometer 




into melting ice, for ice melts at a constant temperature. 

When the position of the mer- 
cury becomes stationary, the 
height of the mercury in the 
stem is marked. The arrange- 
ment for marking the freez- 
ing point' is shown in Fig. 25. 
The funnel is filled with broken 
ice in which the thermometer 
is inserted. Next the whole 
is placed in the steam escaping 
from boiling water in a jacketed 
tin boiler shown in Fig. 26 ; 
and this is also a constant or 
uniform temperature, if the ex- 
periment is always made at 
the same barometric pressure. 
When the height of the mer- 
cury has again become constant, 




Fig. 24. 



LESSON VI THE THREE THERMOMETRIC SCALES 



47 



this point is marked on the stem, 
positions of the freezing and boiling 
points of water (see p. 107 for in- 
fluence of pressure). 

Centigrade, Fahrenheit, and 
Reaumur Scales. — The interval 
between the two points is divided 
differently in different countries, as 
shown in Fig. 27. In the Centi- 
grade (or Celsius) thermometer the 
two points are marked o and 100, 
and the interval divided into 100 
spaces. This scale is the one used 
in this book and in nearly all 
scientific experiments. On the Fah- 
renheit thermometer, which is com- 
monly used in this country, the two 
points are marked 32 and 212 re- 
spectively (the zero point being sup- 



We now have the 





Fig. 25. 



Fig. 26. 



posed by Fahrenheit to 
be the greatest attainable 
cold), and the interval 
divided into 180 spaces, 
each being called one 
degree (1°). In the Reau- 
mur the points are mark- 
ed o and 80, and 80 
spaces or degrees are 
made between them : this 
is used on the continent. 
It is easy to convert any 
temperature from one 
scale to another. If we 
divide the whole interval 
in each by 20, we get 
equal intervals of tem- 
perature which contain 



48 



CONVERSION OF SCALES 



in C. 5, in F. 9, and in R. 4 spaces or degrees, so 5° C. 
F. = 4° R. So by proportion, we can convert • 



212 



100 



Fig. 27. 



14'^ C. into R. 



5 : 4 : : 14 



4 X 14 56 . ^ 
5 5 ' 

180 R. into C. 4 : 5 : : 18 

4 4 

In converting from or into F. scale it 
is rather different. We must always 
take the 7iuviber of degrees above 
freezing point in F. scale before be- 
ginning the conversion from F. to C, 
and we must add 32 to the answer if 
from C. to F., e.g. — 



10° C. into F. 



5:9: 



9 X 10 ^ 
10 = ^ -= 18. 



18 + 32 = 50° F. 



This is because the zero of F. is 32° below freezing point, and 
we always count from zero (or o") ; hence add 32. 

To convert 60^ F. into C. we have only 60-32 = 28^' above 
freezing point, therefore 



What we have learnt 

In our sixth Lesson we have considered the standards of length, area, 
volume, and weight in the metric system, which is now almost universally- 
used for scientific purposes. We have learnt the methods adopted in 
weighing and measuring the volumes of substances by means of the balance 
and graduated glass vessels, such as the flask, pipette and burette. Also 
the conditions of sensitiveness and equilibrium of the balance, and the use 
of the rider in exact weighing. We have learnt something about the expan- 
sion by heat, of solid, liquid, and gaseous bodies, together with the method 
of making and graduating a thermometer, and also the method of convert- 
ing thermometric readings from one scale to another. 



LESSON VI SUMMARY AND EXERCISES 49 



Exercises on Lesson VI 

1. Explain the relations between the standards of length, area, volume, 
and weight in the metric system. How many millimetres are there in 
4.739 metres, and how many centimetres in 3645 millimetres ? 

2. Describe the chemical balance and mention the conditions which 
must be observed to render it sensitive, giving reasons for your answer. 
How would you find out whether the arms of a balance are equal in length, 
and if unequal, how would you find the true weight of a substance by its 
means ? 

3. Describe an experiment to show that glass expands less than water 
for the same increase of temperature. 

4. How are thermometers filled and graduated, and what are the differ- 
ent thermometric scales in use in different countries ? 

5. Convert 4^ C. into F. ; 5^ F. into C. and R. ; 60^ F. into C. and R. ; 18^ 
R. into C. and F. ; -40^ F. into C. 

6. Explain how the rider is used in exact weighing. What is the 
principle upon which it depends ? 

7. What relation exists between the position of the centre of gravity of 
a balance and its sensibility ? 

8. How is the centre of gravity of a balance adjusted ? 

9. Describe an experiment to show that gases expand when heated and 
contract when cooled. 



LESSON VII =^ 
PHYSICAL PROPERTIES OF GASES 

RELATION OF VOLUME TO TEMPERATURE AND PRESSURE 
— DALTON'S LAW — BOYLE'S LAW — CALCULATION OF 
VOLUMES FROM WEIGHTS— REDUCTION TO NORMAL 
TEMPERATURE AND PRESSURE [NTP] 

Relation of the Volume of Gases to Temperature. — 

We have seen in Expt. 20 that gases expand when heated 
and contract when cooled, and also in Expt. 19 that for an 
equal increase of temperature, solid bodies expand less than 
liquids, and these much less than gases ; they also expand 
differently, whilst all gases are found to expand alike or very 
nearly so. 

Law of Dalton.f — Dalton found that when the pressure 
is constant all gases expand very nearly 273- /^^/ of their 
volume at 0° for every increase i7i te7nperature of i^ C. Thus^ 



273 vols. 


of a 


gas 


at 0^ C. become 


274 


a 




i^ 


275 


u 




2° 


276 


u 




3° 


273 + 1 


u 




t° 


ing this we 


can 


convert the volume of any gas 



* In some cases it may be well to postpone this and the following lesson 
until after some of the subsequent lessons have been studied. 

t This is also sometimes called the Law of Charles, but it was first dis- 
covered by John Dalton. 

50 



LESSON VII 



DALTON'S LAW 



51 



measured at a particular temperature to the volume it would 
occupy at any other temperature, e.g. : — 

Calculation of the Volumes of Gases at Different 
Temperatures. — A gas occupies 689 c.c. at 4° C, what vol. 
will It occupy at 15° C? 273 + 4 or 277 vols, at 4° C. become 
273+ 15 or 288 vols, at 15°, and therefore 

„„ ^„ 288 X 689 

277 : 288 : : 689 = = Ans. 

277 

Directions. —Add 273 to each temperature. Then, as the third term 
of a proportion sum place the given volume, and in the second term place 
the higher number if the gas expands by heating from a lower to a higher 
temperature ; and the lower number if the gas contracts by cooling from a 
higher to a lower temperature, e.^. : — 

Second Method. —These calculations also maybe worked out by the 
method of reduction to unity explained on p. 35. 

A gas occupies 1000 c.c. at 20^ C, what vol. will it occupy 

ato^C? 

273 + 20 = 293 

1000 X 273 

293 : 273 : : 1000 = — = Ans. 

^^ ^^ 293 

ExPT. 21. Relation of the Volume of Gases to 
Pressure. — If we take a strong glass tube about 3 feet long, 

closed at one end, and fill it with ,.; 

liquid mercury, and then clos- i j """---..^ 

ing up the end tightly with the 
thumb, invert it in a basin full 
of mercury, we shall see on re- 
moving the thumb, that the mer- 
cury falls in the tube and leaves 
a space in the upper part, of 
about 6 inches long. If we 
incline this tube as shown in 
Fig. 28 to the second position 
B, in which the end of the tube 
falls below the level of the mer- 
cury in the first position A, we 
see that the mercury rises and 
completely fills the tube, show- 
ing that the space C of 6 inches 




>3olN. 



Fig. 28. 



did not contain any gas. This experiment was first made by 



52 



EXPERIMENTAL PROOF 



Torricelli, and the space C is called a Torricellian vacuum. 
Why does the mercury fall? and why does it rise in the tube 
when it is inchned? Why does not all the mercury run out? 
What keeps it up? These are the questions we may ask our- 
selves. The column of mercury is 
held up by the pressure of the atmo- 
sphere. The atmosphere rises to a 
height of at least 40 miles above the 
surface of the earth, and the accumu- 
lated weight of the mass presses on 
the mercury in the basin, and this in 
turn presses on the mercury in the 
tube, upholding a column of about 
30 inches in height. When the tube 
is inclined as at B, the vertical height 
is less than 30 inches, so the mer- 
cury completely fills the tube. This 
column of mercury is the essential 
part of a mercurial bay^ometei' or pres- 
sure measurer. With atmospheric 
changes the pressure varies and the 
height of the column of mercury 
varies between wide limits, 2 or 3 
inches frequently. 

Let us now try some experiments 
to find out the relation between the 
volume of gases and the pressure to 
which they are subjected. 

ExPT. 22. Boyle's La'w. — Into 
the funnel A (Fig. 29) w^e will pour 
mercury until it is at level B in both 
tubes (the stopcock C being open). 
Closing C w^e enclose a unit volume of 
air at the ordinary pressure of the 
atmosphere, say the barometer is 
standing at 30 inches, let us now 
pour mercury into A until the difference in the levels in the 
[J tube is 30 inches. Then in addition to the ordinary atmo- 
spheric pressure w^e have added a pressure of mercury equal to 
another atmosphere ; so our air is under a pressure of tw^o 




Fig. 29. 



LESSON VII OF BOYLE'S LAW 53 

atmospheres, and we find it has diminished or become com- 
pressed into just hah' its original voUime. Adding more mer- 
cury until the difference in level is 60 inches, thus subjecting the 
air to 3 atmospheres' pressure, we find the volume reduced to J ; 
whilst at 90 inches or 4 atmospheres the volume is i- 

Thus, pressures i, 2, 3, 4 (atmospheres) 
gives volumes i, J, J, i, 

or the volume is inversely proportional to the pressure. 

Gradually letting out mercury at D, we notice the volume of 
air gradually increases. Beginning at 90 inches with } vol. at 60 
inches we get ^, at 30 inches ^, at 15 inches f , and at 10 inches 
5, whilst at the level at B (if our stopcock C has been perfectly 
tight) we again obtain the original volume of air ; in diminishing 
the pressures the volumes were : — 

Volumes J, 4, i, f, J, i, 

Pressures 4, 3, 2, i, |, i (atmospheres). 

Exactly the same results would have followed if instead of air we 
had taken hydrogen, oxygen, or nitrogen. 

This experiment illustrates Boyle's Law, which states that 
when the temperatitre is constant the voliune of a gas is iiiversely 
p7'oportional to the pressure. But it only proves it for pressures 
greater than that of the atmosphere. 

It also follows from Boyle's Law and from the above experi- 
ments that the density of a gas is directly proportional to the 
pressure to which it is subjected. 

ExPT. 23. — To prove the law for lower pressures than that 
of the atmosphere, the apparatus shown in Fig. 30 must be 
employed. It consists of a long tube full of mercury, at the 
top of which is a glass dish, also containing mercury. A tube, 
closed at one end, and about 5 feet long, is filled with mercury 
except about 6 inches, the end is firmly closed with the thumb 
and inverted in the glass mercury trough. The tube is now 
depressed in the trough until the mercury is at the same level 
inside and outside the tube, the contained air is then at the 
ordinary pressure, say 30 inches, a little strip of paper is 
gummed to mark the volume of air, and the tube raised until 
the volume is doubled ; it will then be seen that the mercury 



54 



CALCULATION OF VOLUMES 



PART 



has risen 15 inches. The total pressure at the level of the 
mercury in the trouo^h is 30 inches, and as we have mer- 
cury in the tube equal to 15, the gas 
must press with a force equal to an- 
other 15, making 30, or the pressure 
of the gas is found by subtracting the 
height of the mercury column (meas- 
ured from the level in the trough), 
from the height of the barometer ob- 
served at the time. 30 — 15 = 15 
or i an atmosphere. If it is now 
raised until we have 3 times the vol- 
ume, we find our column at 20 inches, 
so pressure is 30 — 20 = 10 or | an 
atmosphere. Similarly, when we 
have 4 volumes the pressure is 30 — 
22 J = 7J = J atmosphere. Again we 
have 



Pressures i, ^, |, J (atmospheres), 
Volumes i, 2, 3, 4. 

Boyle's Law is not correct for any 
gas under very high pressures nor in 
the case of gases which are near their 
point of liquefaction (see p. 63). 

Calculation of the Volumes of 
Gases at different Pressures. — 
Suppose it is desired to find the vol- 
ume of a gas at a pressure of 785 mm. 
which at 735 mm. measures 478 c c. 
From Boyle's Law we know that the 
volumes are inversely proportional to 
the pressures, and, therefore, in the 

ratio of 785 : 735. So the volume will be — — ^ cc. It is 

easily remembered whether to multiply by the higher or lower 
number, by asking, Is the volume increased or diminished ? If 
the pressure is diminished the volume is increased, and we must 
multiply by the higher and divide by the lower number and 
vzce versa. 



LESSON VII GAY-LUSSAC'S LAW 55 

Calculation of the Volumes of Gases at different 
Temperatures and Pressures. — Making both corrections 
at the same time, we may wish to find, for example, the volume 
of hydrogen at NTP (normal temperature and pressure, o° C. 
and 760 mm.), which at 15° and 780 mm. measures 587 c.c. 

587 X 273 ^ 780 
288 760' ^ 

587 X 273 X 780 



288 X 760 



= vol. of hydrogen at NTP. 



It will be noticed that the gas is cooled, and its volume 
therefore tends to diminish, and that the pressure is lowered, 
and therefore its volume tends to increase. 

Relations between the Densities of Gases and their 
Atomic and Molecular Weights. — Gay-Lussac found, tak- 
ing the density of hydrogen as i, that t/ie numbers expressing 
the densities of the gaseous elements are identical with their 
atomic w eight s^"^ whilst the deiisity of any compound gas is half 
its molecular weight. We shall afterwards find that the molecule 
of the elementary gases consists of two atoms. The law may, 
therefore, be stated as follows: the density (H = i) ^/" ariy gas 
{^elementary or co7npound) is half its molecular weight. 



Elementary Gases. Atomic Weight. 


Molecular Weight. 


Density. 


Hydrogen 


I.O 


2.0 


1.0 


Oxygen 


iS-9 


31.8 


iS-9 


Nitrogen 


13.9 


27.8 


13-9 


Chlorine 


35.2 


70.4 


35-2 


Compound Gases. 


Formula. 


Molecular Weight 


Density. 


Water 


H2O 


17.9 


8.95 


Nitrogen Monoxide 


N2O 


43-7 


21.85 


Carbon Monoxide 


CO 


27.8 


13.90 


Carbon Dioxide 


CO, 


43.7 


21.85 



Calculation of Volumes of Gases at any Temperature 
and Pressure, from a given Weight of Gas. — We have 
already learnt that hydrogen is the lightest gas known, and 

* There are certain exceptions to this law which will be noticed later, 
e.g. phosphorus and arsenic have 4 atoms to the molecule, mercury has 
only one, and the density of phosphorus and arsenic is double the atomic 
weight, whilst that ot mercury is only half, 



56 CALCULATION OF VOLUMES part I 

that it is taken as the standard ^Yith which to compare the 
densities of all the other gases. It has been "found by careful 
experiment that a litre of Jiydrogen at N'T P weighs 0.0899 g^'^^^'i- 
It is evident that the weight of a litre of any gas is proportional 
to its density, and therefore a litre of oxygen at NTP weighs 
0.0899 grams x 15.9, and a litre of nitrogen at NTP weighs 
0.0899 ^ 13-9 because oxygen is 15.9 times, and nitrogen 13.9 
times as heavy as hydrogen, whilst a litre of carbonic acid 
gas at NTP weighs 0.0899x21.85 (grms.). If a litre of 
hydrogen weighs 0.0899 g^'^^^'s., it is easily calculated that 
2 grms. of hydrogen occupy 22.247 litres (say 22.25 ^^ 22} 
litres), therefore 

2 grams (or 1x2) of hydrogen occupy 22.25 litres 

31.8 " (15.9 X 2) " oxygen " 22.25 " 

27.8 " (13.9 ^ 2) " nitrogen " 22.25 

70.4 " (35.2 X 2) " chlorine " 22.25 

43.7 " (21.85 ^ 2) " carbon dioxide " 22.25 " 

In other w^ords. the molecular weight (in grams) of any gas 
occupies 22.25 litres. 

Examples. — (i) What volume of hydrogen at 15° and 
740 mm. will be given off by adding 4 grms. of sodium to 
water ? 

H.,0+ Na^NaOH + H. 



Directions. — First Method. 

(i) Find the weight {iv) of hydrogen given off (see p. 34). 

(2) Find the volume (^') at XTP corresponding to w. 

-- volume in cubic centimetres (XTP) = v . 

(3) Find the volume corrected to the required temperature and pressure. 
?7 X288 X 760 



I X 4_ 
22.9 

W X lOOO 



273 X 740 



: V the vol. sought. 



Second Method. — A shorter method than the foregoing may be 
carried out as follows, depending upon the fact that the molecular weight of 
a gas (in grams) occupies 22.25 litres ; or, the atomic weight of an elemen- 
tary gas, such as hydrogen, oxygen, chlorine, occupies 11. 125 litres. 

H2O + Na = NaOH + H. 

22.9 grams. xx.\2.^ litre- 



LESSON VII OF GASES FROM THEIR WEIGHTS 57 

In the foregoing equation we may take it, therefore, that 22.9 grams 
of sodium yield 11. 125 litres of hydrogen (NTP). Then by reduction to 

unity; if 22.9 yield 11. 125 htres, i yields H£^, and 4 yields, therefore, 

22.9 

II. 125 X 4 ^ Qj. 44-5 litres at NTP, this volume must then be corrected to 

22.9 22.9 

15^, and 740 mm. by the method already given. 

(2) What volume of carbon dioxide CO2 at 5^ and 680 mm. 
will be given off by burning 15 grams of carbon? C + 02 = C02. 

43.7 X 15 wx 1000 ty X 278 X 760 ^_ ,. , , 

^^-^ -=iv — 7^ = y — 77~ — -=V(mc.c.s). 

11.9 .0899X21.85 273x680 

C + O2 = CO2. 

II grams. 22.25 litres. 

The above equation may be read that 11.9 grams of carbon yield 22.25 

litres of CO2 at NTP, then we have by reduction to unity ?^5J<Jt5^ x\iQ 

11.9 
required volume at NTP, which must be corrected to 5^ and 680 mm. 

Calculation of Weights from Volumes of Gas at any 
Temperature and Pressure. — It must always be remembered 
that dividing the weight of any gas by 0.0899 ^ ^ (where d is 
its density) gives its volume at NTP, and this volume must then 
be corrected for any other temperature and pressure. On the 
other hand, when it is required to find the weight of any volume 
of gas at any temperature and pressure other than NTP, its vol. 
must first be converted to NTP before proceedmg to find the 
weight, e.g.: — 

(3) What weight of carbon is contained in 4.78 litres of carbon 
dioxide measured at 15° C. and 660 mm. ? 

Directions. — First Method. 

(i) Find vol. at NTP. 4-78X273x660 ^^^ .^^^^ ^^ ^ 

^ ' 288 X 760 ^ ^ 

(2) Find weight of gas. v X .0899 X 21.85 = "^ (i" grams). 

1 1 O X 1i) 

(3) Find weight of carbon. — ~ = W, weight of carbon. 

Second Method.— After reducing to NTP, we may proceed thus: 



58 SUMMARY AND EXERCISES part I 

22.25 litres of CO2 (NTP) correspond to 11.9 grams of carbon, then by 
reduction to unity we have 

22.25 litres CO2 equal 11.9 grams of C. 

„ El^ .. 

22.25 

y " «' Z " 

22.25 

where v is the volume of CO2 in litres reduced to NTP. 

What we have learnt 

In our seventh Lesson we have considered Dalton's Law, which states 
that all gases expand very nearly 073 part of their volume at o*^ for every 
increase in temperature of 1° C, and also the method of converting gaseous 
volumes from one temperature to another. The relations of the volume 
of gases to pressure has been studied; it was seen that all changes in 
volume and pressure conform to Boyle's Law, which states that the volume 
of any gas is inversely proportional to the pressure to which it is subjected. 
We have seen how the law may be proved for pressures both above and 
below that of the atmosphere, and also how to calculate gaseous volumes 
from one temperature and pressure to any other. The relations between 
the atomic and molecular weights of gases, and their densities have been 
considered, as also has the method of calculating the volumes of gases 
from their weights, and vice versa. 



Exercises on Lesson VII 

1. What is Dalton's Law ? Convert 869 vols, of gas at 14° C. to 5° C. 

2. What is a Torricellian vacuum ? How can you prove it does not 
contain any gas ? 

3. How would you demonstrate Boyle's Law both for pressures above 
and below that of the atmosphere ? 

4. Convert {a) zyS c.c. of gas at 3° C. and 640 mm. to 16° C. and 780 
mm. ; (d) 1000 c.c. of gas at o*^ and 760 mm. to 15^ C. and 720 mm. 

5. What weight of oxygen is contained in 10 litres of carbon dioxide 
measured at 12= C. and 770 mm. ? 

6. What volume of carbon dioxide at 20° C. and 725 mm. will be given 
off from 560 grams of marble, CaCOs wdiich is completely decomposed by 
heat into lime, CaO and carbon dioxide, CO2 ? 

7. Calculate the density of the following gaseous bodies : ozone, O3, 
hydrochloric acid, HCl, sulphuretted hydrogen, H2S, ammonia, NH3. 
What is the weight of 10 litres of ammonia gas at NTP ? 

8. What weight of sulphur must I burn in order to generate 60 litres 
of sulphur dioxide, SO2, at 20^ and 760 mm. ? 



LESSON VIII 

PHYSICAL PROPERTIES OF GASE^— Continued 

RELATIONS BETWEEN THE COMBINING VOLUMES OF 
GASES — AVOGADRO'S LAW — DIFFUSI ON, LIQUEFAC- 
TION, AND SOLIDIFICATION OF GASES 

Avogadro's Law. — Since the densities of the elementary 
gases are indicated by their atomic weights (p. 55), it follows 
that eqtial vohimes of any of the eleiJientary gases co7itain the 
same number of atoms at the saine temperatitre a7td pressure. 
This law may be extended to include all gases, since the den- 
sity of any gas is half its molecular weight, and we may say 
therefore, that equal volumes of any of the gases, ele7ne7ttary or 
co?npound, contain the sa?ne nu7nber of molecules * (at the sa77ie 
te77iperature and pressure') . This is known as Avogadro's Law. 

Relations between the combining Volumes of Gases. 
— It follows, from Avogadro's Law, that the ratio between 
the combining volumes of gases must be a very simple one. 
In this connection it is always to be taken that the atoin of 
hydroge7i occupies one volimie, and the molecule or two atoms 
of hydrogen occupies two volumes. Since equal volumes of 
different gases contain the same number of molecules it follows 
that the molecule of any gas simple or compound, occupies the 
same volume as a molecule of hydrogen. In other words the 
7fiolecule of any gas occupies two vohmies. This law, it must 

* The molecule of a compound is the smallest group of its component 
atoms which can exist in the free state, hence the molecule of a compound 
cannot be divided without decomposing it. 

59 



6o COMBINATION BY VOLUxME part I 

be remembered, is deduced from Avogadro's Law, which is 
founded upon the fact, discovered by Gay-Lussac, that the 
densities of the elementary gases are denoted by the numbers 
expressing their atomic weights. 

Taking a few examples of the simple relations between the 
combining volumes of gases, it has been found by actual 
experiments, which will be considered under their respective 
headings, that : — 

(i) When four volumes of hydrogen and two volumes of oxygen 
combine together to form water vapour as steain, a contraction 
takes place and/<?//r volumes of steam are produced (see p. 94). 

2H2 + O, = 2H2O. 

2 Molecules. Molecule. 2 Molecules. 
4 vols. 2 vols. 4 vols. 

(2) When carbon or sulphur are burnt in oxygen in a 
closed vessel no increase in the volume of the gas is observed 
(see p. 203). 

C + O. = 00, S + O, = SO2. 

(Solid). Molecule. Molecule. (Solid). Molecule. Molecule. 
2 vols. 2 vols. 2 vols. 2 vols. 

(3) When nitrogen monoxide N^O is decomposed by means 
of potassium, w^hich combines with and absorbs the oxygen 
and sets the nitrogen free, no change of volume takes place 
(see p. 155). 

N,0 + K2 = No + K2O. 

Molecule. (Solid). Molecule. (Solid). 

2 vols. 2 vols. 

(4) When ammonia gas (NHo) is decomposed by means of 
the electric spark, its volume is doubled (see p. i66). 



2 NH3 


= N, 


+ 3 Ho. 


2 Molecules. 


Molecule. 


3 Molecules. 


4 vols. 


2 vols. 


6 vols. = 8 vols. 



In all these cases we find that the deductions which we have 
made from theory, are confirmed by actual facts obtained by 
experiment. 

Diffusion of Gases. — Dobereiner first noticed in 1823 
that when a flask which happened to have a fine crack in it, 



LESSON VIII 



DIFFUSION OF GASES 



6i 



was filled with hydrogen over water, and then allowed to stand, 
the volume of the hydrogen was found next day to have become 
much less. The same diminution did not take place if the 
cracked flask was covered with a bell-jar containing hydrogen, 
nor did it take place if the flask was filled wdth air. Graham, 
in 1832, first showed that air entered the flask when hydrogen 
left it, but in the first case, hydrogen passed out or diffused 
through the ox^iok faster than air entered, whilst in the second 
case hydrogen was able to enter or diffuse into the cracked 
vessel just as fast as it diffused out, 
and no alteration of bulk w^as to be 
noticed. 

ExPT. 24. — This power of the mole- 
cules of gases to diffuse through very fine 
apertures is best illustrated by means of 
the following experiment. A cyhndrical 
porous cell, such as is used for voltaic 
batteries, is fitted with an india-rubber 
stopper pierced with a hole into w^hich a 
long glass tube is fixed. The tube has a 
bulb and U tube ending in a jet (Fig. 31), 
and is filled with a coloured liquid before 
the experiment, whilst the porous cell 
contains air. If now a beaker full of 
hydrogen is brought over the porous cyl- 
inder, it will be found that the coloured 
liquid is thrown up as a fine fountain 
showing that the pressure in the porous 
cell is increased. The reason is, that 
the hydrogen diffuses or passes through 
the pores of the porous cell 77iiich faster than the air diffuses out 
of it. 

ExPT. 25. — If wx arrange the apparatus rather differently 
(as in Fig. 32), so as to surround the porous cell (containing 
air), with the heavy gas carbon dioxide, wt shall find that the 
air passes out of the porous cell more rapidly than the carbon 
dioxide diffuses into it ; the pressure is thereby diminished and 
air passes from the inlet tube, through the water and into the 
porous cell to supply the deficiency. 

ExPT. 26. — If w^e take two small gas jars with ground and 




Fig. 31. 



62 



DIFFUSION OF GASES 



ii 



well-fitting mouths, place them together (Fig. 33), and render 
them air-tight by rubbing vaseline over the ground surfaces, 
we might expect, at first sight, that if 
the upper jar is completely filled with 
hydrogen (density =1), and the lower 
one with carbon dioxide 
(density = 22), two gases 
which do not combine with 
each other, then the heavy 
gas would reiJiain in the 
lower jar and the light gas 
in the upper one. But this 
is not found to be the case. 
Our previous experiments 
on diffusion would prepare 
us for the observed fact 
that, in two co7n7nunicating 
vessels containing different 
gases, no matter what their 
relative densities, diffusion 
takes place until an equili- 
brium is established, and then both 
vessels contain an identical mixture of 
the two gases. If we try the experiment, and leave the two 
gases in contact as shown in Fig. 33, for a few hours, then 
slipping a circular glass plate over the mouth of each jar, 
invert them over a strong solution of caustic potash, we shall 
find that the liquid rises half-way up each jar, showing that 
both jars contain an equal mixture of carbonic acid gas (which 
is absorbed by caustic potash), and hydrogen (which remains 
unabsorbed). 

ExPT. 27. — Another experiment, showing the rapid diffusion 
of gases, may be easily made as follows. A gas jar is 
inverted in a retort stand ring, and filled with hydrogen 
by upward displaceirient (see p. 72). Another jar (not 
inverted), is filled with carbonic acid gas by downward dis- 
placement (see p. 225). The two jars are allowed to remain 
open in a roomj of equable temperature and free from draughts 
for a few hours. When a lighted taper is introduced into 
both jars it continues to burn in each as in air. In fact both 




Fig. 33. 



Fig. 32. 



LESSON VIII LIQUEFACTION OB^ GASES 63 

jars are now full of ordinary air. Both the hydrogen and the 
carbonic acid have diffused into the atmosphere. 

Law of Diffusion. — Graham found that the relative rates of 
diffusion of gases are inversely proportioiial to the square roots 
of their densities. 

Hence hydrogen will diffuse four times as fast as oxygen, 

for -^: — ^ : : 4: I. In Expt. 27 if we had ascertained the 

Vi V16 
times taken for the gases to diffuse from the jars, we should 
have found that the hydrogen diffused away much more rapidly 
than the carbon dioxide. The following table gives the 
rate of diffusion of several gases, as experimentally determined 
by Graham, compared with the inverse square roots of their 
densities. 

J Velocity of 

Density. - Diffusion. 

Air=i. ^Density. Air = i. 

Hydrogen 0.06926 3-7790 3-830 

Nitrogen 0.97130 1.0150 1.014 

Oxygen 1. 10560 0.9510 0.949 

Carbon dioxide 1.52900 0.8087 0.812 

showing that the numbers found by experiment agree closely 
with those calculated from the law. 

The diffusion of gases is of very great importance in the 
economy of nature, for by this means the air of towns and 
dwellings is kept in a pure condition, the vitiated and foul air 
constantly diffusing into large volumes of purer air. Diffusion 
however is only secondary to the action of winds and ventila- 
tion,* in the renewal of the pure air which we require for 
breathing. 

Liquefaction of Gases. — We have seen in our first les- 
son that all gases may be condensed to liquids by the 
combined influence of cold and pressure. Some substances 
which are gases at the ordinary temperature may be con- 
densed to liquids at the ordinary atmospheric pressure by 
means of cold alone, whilst at o^ these gases may be con- 
densed to liquids by pressure alone as shown in the following 
table: — 

* See p. 145 for an account of ventilation. 



64 CRr 


nCAL POINT 




Part i 




Cold A one. 


Pressure Alone. 




(Pressure 760 mm.) 


(Temperature, 0° C.) 


Sulphur dioxide SO2 


— 10° 


1.53 


Atmospheres 


Chlorine CI 


-34° 


6 




Sulphuretted hydrogen H2S 


— 62^ 


10.8 




Ammonia NH3 


-40 


4.40 




Carbon dioxide CO2 


— 80 


38.50 




Nitrous oxide N2O 


-92 


32 





Critical point in the Liquefaction of Gases. — The case of 
carbon dioxide is very' instructive, because it' is found impossible 
to liquefy that gas by pressure alone if the temperature is higher 
than a certain limit. Above this temperature, if the gas be com- 
pressed until a pressure of, say, 20 atmospheres is reached, no 
liquefaction takes place, but if now the tube containing the com- 
pressed gas be gradually cooled, we find that at the temperature 
30.92^ the gas becomes disturbed and cloudy and separates into 
a layer of liquid carbonic acid, which is sharply distinguished 
from the gas above. If now the temperature be raised above 
30.92°, the liquid appears to lose its sharp surface and the tube 
is filled with a cloudy mass in commotion and soon nothing 
is visible, the liquid having become entirely vaporised. This 
temperature (30.92°) is called the critical teinperattire of lique- 
faction for carbon dioxide. Above it the gas cannot be liquefied, 
no matter how great the pressure. There are similar but not 
identical critical temperatures for all gases. The so-called per- 
77ianent gases such as oxygen, hydrogen, nitrogen, were for a 
long time considered to be non-liquefiable, simply because 
pressure was brought to bear upon them at temperatures 
higher than the critical point. We now know that their criti- 
cal points are far below the freezing point (o^), and by exces- 
sive cold as well as great pressure they have been obtained as 
liquids. 

Heat produced when Gases are compressed. — This is 
very well shown by means of a strong glass tube fitted with 
a light piston. When the tube is full of air and the piston 
forced inwards so as to compress the air, heat is evolved which 
is sufficient to inflame a small piece of tinder moistened 
with ether, which is inserted into a cavity at the end of the 
piston. 

Cold produced when Gases expand. — Conversely when 
gases expand, heat is absorbed or cold produced. This 



LESSON VIII SUMMARY AND EXERCISES 65 

absorption of heat is sufficient to cause the soUdification of 
carbonic acid gas when the Hquid is run out from a strong iron 
bottle through a fine jet into a brass box. Part of the liquid 
vaporises at once, and the heat absorbed in its conversion from 
liquid to gas is so great as to solidify the remaining liquid. 
Solid carbon dioxide is a snow-like mass and is often used 
for obtaining very low temperatures. 

Mixed with ether and placed in the vacuum of an air pump 
its evaporation produces a temperature of — 100° C. 

What we have learnt 

In our eighth Lesson we have considered Avogadro's Law which states 
that equal volumes of any of the gases contain, at the same temperature 
and pressure, the same number of molecules. 

We have seen that a very simple relation exists between the combining 
volume of gases, and that the molecule of any gas occupies the same 
space (2 volumes) as the molecule (2 atoms) of hydrogen. 

The diffusion of gases has been demonstrated and the law of diffusion 
discovered by Graham has been given. We have learnt something more 
about the liquefaction and solidification of gases, and noticed the influence 
of temperature near the critical point. 



Exercises on Lesson VIII 

1. What is Avogadro's Law ? From what experimental facts is it 
deduced ? 

2. What can be deduced from Avogadro's Law as regards the relations 
between the combining volumes of gases ? 

3. What volume of chlorine will be required to completely decompose 
10 litres of sulphuretted hydrogen, H.2S ? [CI.2+ H.2S = 2 HClV S (solid)] , 
and what volume of hydrochloric acid gas, HCl, will be formed ? 

4. How would you demonstrate the diffusion of hydrogen and carbon 
dioxide through the walls of a porous cell containing air ? 

5. What is meant by the critical point in the liquefaction of gases ? 
Give an example. Why were nitrogen and oxygen considered to be 
perinanent gases ? 

6. How may carbon dioxide be solidified ? 

7. State Graham's law of the diffusion of gases. 

8. How was the diffusion ' f gases discovered ? 

9. Give examples of the snnple relations existing between the combin- 
ing volumes of gases. 



PART II 

SYSTEMATIC STUDY OF CERTAIN NON- 
METALLIC ELEMENTS, AND THEIR 
MORE IMPORTANT COMPOUNDS 



LESSON IX 
Hydrogen 

Symbol H. Atomic Weight i. Density i 

Cavendish in 1766 first ascertained the true nature of this 
gas to which he gave the name of inflammable air. 

Occurrence. — Hydrogen occurs almost solely in a state of 
combination in nature, though it sometimes exists in the free 
state mixed with other gases in certain volcanic emanations. 
Its principal compound is water, HgO, of which it forms one- 
ninth part by weight. As all the oceans, seas, rivers, and 
lakes on the face of the earth contain one-ninth their weight 
of hydrogen, the quantity of this element occurring in a state 
of combination is very large. Hydrogen also forms an essential 
ingredient of all animal and vegetable bodies, and most organic 
compounds {e.s^. sugar, starch, fat, wax, etc.) contain it as a 
constituent. All mineral oils contain a large proportion of 
hydrogen combined with carbon. 

67 



68 



PREPARATION OF HYDROGEN 



Test for Hydrogen. — Hydrogen may be recognised by 
the fact that if brought in contact with a flame, it burns with 
a pale lambent blue flame, and, if burnt in a dry jar, bedews 
the sides of the vessel with moisture, owing to its combination 
with oxygen (in the air) to form water, whilst if lime-water be 
subsequently added, no precipitate is produced as would be the 
case if the gas had been a carbon compound, some of which 
burn similarly. 

Preparation. — (i) The usual method of preparing hydro- 
gen is by the action of dilute sulphuric or hydrochloric acid on 
metallic zinc, the chemical reactions are — 



Zn + H2SO4 = 

Zinc. Sulphuric acid. Zinc sulphate. Hydrogen. 



ZnSO^ + H2. 



Zn + 2HCI = ZnClg + Hg. 

Zinc. Hydrochloric acid. Zinc chloride. Hydrogen. 



ExPT. 28. — Into the flask (Fig. 34), a quantity of granulated 
zinc is introduced. The flask is fitted with an india-rubber 
stopper pierced with two holes. Into one of these is fitted a 
long thistle funnel, reaching nearly to the bottom of the flask, 
and into the other a bent delivery tube passes just beneath 
the stopper, and thence beneath the bee-hive shelf of the 
pneumatic trough, and serves for the delivery of the gas into 

the gas jar which is first 
filled with water. Having 
arranged the apparatus, 
the dilute acid is poured 
down the thistle funnel so 
as to cover the zinc to a 
depth of one or two inches. 
The liquid soon begins to 
effervesce, from the evolu- 
g tion of hydrogen gas. Here 
a most important caution 
must be given to prevent 
accidents. Beginners, when 
making this experiment, must remember that hydrogen gas and 
common air for7ii a highly explosive mixture', and that, as 
the flask is filled with air, the gas which first comes off is a 




Fig. 34. 



LESSON IX 



PREPARATION OF HYDROGEN 



69 



mixture, which, if collected in a gas jar and a light applied, 
will explode violently. Therefore we must, before beginning 
to make any experiments zvith hydrogen^ be quite sure that no 
air is mixed with it. To ensure this we have only to wait 
several minutes until the hydrogen, which is quickly evolved, 
has driven out the air. We can easily learn whether this has 
taken place, and whether the gas is ready for collecting, by 
placing a test-tube filled with water, and inverted, over the 
little hole in the bee-hive shelf. When the tube is full of 
gas it must be removed, with the thumb closing its mouth, 
and held mouth downwards to a flame. If the gas burns with 
a slight explosion, or does not burn at all, the gas is stiJl 




Fig- 35- 

mixed with air, and we must wait a little longer; and if after a 
short time the trial is again made, and the gas burns quietly, 
the gas is free from air and may be collected for further ex- 
periments (Experiment 44, p. 96, shows the explosive nature of 
a mixture of hydrogen and oxygen). 

(2) Iron filings may be used instead of zinc, but the gas thus 
evolved is not so pure owing to impurities in the iron. The 
reaction is similar to that which occurs when zinc is used. 



Fe + H2SO, = FeS04 -f H2. 

Iron. Sulphuric acid. Ferrous sulphate. Hydrogen. 

Fe + 2HCI = FeCl, + Hg. 

Iron. Hydrochloric acid. Ferrous chloride. Hydrogen, 



70 



ACTION OF STEAM ON RED-HOT IRON part 



(3) By the action of sodium * on water (Fig. 35 ; see also 
Experiment 11, p. 14), 

2H2O + Na^ = 2NaOH + Hg. 

Water. Sodium, Caustic Soda. Hydrogen. • 

The hydrogen obtained in this manner generally burns with 
a bright yellow flame owing to the presence of traces of 
sodium hydroxide. 

(4) By the action of water as steam on red-hot iron, 

3Fe + 4H2O = FcgO^ + 4H2. 

Steam. Ferrosoferric oxide. 

Water is boiled in a small flask, and the steam passes into 
an iron or porcelain tube filled with iron borings heated to 




Fig. 36. 

redness in the furnace (Fig. 36). The oxygen of the water 
(steam) combines with the iron to form the black oxide of 
iron, Fe304 (magnetic oxide of iron, or ferrosoferric oxide), 
whilst the liberated hydrogen passes on into a gas jar. 

(5) By the electrolysis of acidulated f water (Fig. ^7 ; see 
also Experiment 12, p. 15), 

2 H^O r= 2 H, + O2. 

* Potassium also decomposes water in a similar manner, but the action 
is more violent and the heat evolved is sufficient to set fire to the gas, 
which burns with a violent flame owing to the presence of potassium 
hydroxide. 

fA little sulphuric acid is added to enable the water to conduct the 
current of electricity. 



LESSON IX 



PROPERTIES OF HYDROGEN 



71 



Properties. — Hydrogen is a colourless, invisible, tasteless, 
and inodorous gas, and is the lightest substance known, and 
for this reason it is taken as 



the standard with which to 
compare the density or heavi- 
ness of all the other gases. 
Hence at the head of this 
lesson we find Density = i, 
or the density of hydrogen is 
taken as unity. The density 
of common air at NTP is 
14.39, ^vhilst the density of 
carbonic acid gas is 21.85. 
In other words air is nearly 
14 J and carbonic acid nearly 
22 times as heavy as hydro- 
gen. A litre of hydrogen at 
NTP weighs 0.0899 grams. 




Fig. 37- 



This number should always be remembered, because, knowing 
the density of any gas, the weight of a litre of it is obtained by 
multiplying 0.0899 by the density. 



Experiments -with Hydrogen 

ExPT. 29. Hydrogen a very light Gas. — (i) Collect a 
jar of the gas and pour it upwards 
into a similar inverted dry gas jar 
(Fig. 38 ; see also Experiments 2 and 
3, p. I . The light gas there mentioned 
is hydrogen). The light hydrogen as- 
cends to the top of the jar and drives 
out the air before it as it gradually 
fills the jar. A lighted taper brought 
to the mouth of the jar will ignite the 
gas with a slight explosion, because 
in pouring the gas upwards it has be- 
come mixed with a little air. Notice 
that the previously dry sides of the 
gas jar become bedewed with mois- 
ture, showing that when hydrogen burns in air it produces water. 




Fig. 38. 



72 



UPWARD DISPLACEMENT 



ExPT. 30. Hydrogen may be Collected by Up-ward 
Displacement. — (2) The bent delivery tube in the generating 
apparatus is replaced by an upright one passing to the top of an 
inverted gas jar (Fig. 39). The light gas collects at the top 
and displaces the heavier air. That the jar is full of hydrogen 
may be shown by lighting the gas at the mouth of the jar. If 
after collecting the gas by upward displacement it be lighted 

at the jet, and a cold bell jar 




I^'ig- 39- 



be held over the flame, it will 
be seen that the inside of 
the bell jar becomes bedewed 
with moisture, which gradu- 
ally collects into small drops 
of liquid water. 

ExPT. 31. Soap Bub- 
bles blown with Hydro- 
gen ascend. — (3) The gas 
is passed through a (J tube 
filled with cotton wool, and 
the delivery tube is connected 
by means of a piece of india- 
rubber tubing to a thistle 
funnel, which is placed in a 



soap solution * so as to blow bubbles On breaking away they 
ascend in the air and rise to the ceiling where they burst. 
Light collodion balloons when filled with hydrogen, rise in the air. 
ExPT. 32. Hydrogen is an Inflammable Gas but does 
not support Combustion. — (4) Into an inverted jar of 
hydrogen plunge a small lighted taper fastened on the end of 
a piece of wire. The flame will set fire to the hydrogen at the 
mouth of the jar, but will itself be extinguished when raised into 
the gas. On withdrawing the taper it again takes fire on com- 
ing in contact with the burning gas near the mouth ; plunged 
upwards it again goies out, and may be extinguished and re- 



* The soap solution for these experiments is best prepared by dissolving 
10 grams of sodium oleate in 400 c.c. of cold distilled water ; 100 c.c. of 
pure glycerine are then added, and the mixture well shaken. After allow- 
ing to stand in the dark for a few days, the clear solution is siphoned off, 
and a drop of strong ammonia added. It should be kept in the dark and 
not exposed to the air. 



LESSON IX SUMMARY AND EXERCISES 73 

lighted many times by moving it into and out of the gas. This 
shows that both the hydrogen and the taper only burn when 
in contact with the air. Hydrogen cannot directly combine 
with the materials of the taper (carbon and hydrogen), and 
the flame is, therefore, extinguished in that gas. 

What we have learnt 

In our ninth Lesson we have learnt the principal methods of preparing 
hydrogen and the most important of its properties. It may be made, e.g. 
(i) by the action of sulphuric or hydrochloric acid on zinc or iron; (2) 
from water, {a) by electrolysis, {h) by the action of sodium or potassium, 
or {c) by the action of red-hot iron on steam. Its principal properties 
were seen to be its lightness and inflammability, and its want of power to 
support combustion in the ordinary manner like air. Hydrogen is a 
colourless, invisible, inodorous gas. Gaseous water or steam is produced 
when hydrogen burns in air. Hydrogen can be collected by the upward 
displacement of air. 

Exercises on Lesson IX 

1. Describe two methods of preparing hydrogen, give equations for the 
reactions, and sketch the apparatus you would use. 

2. What precautions must you take in showing experiments with 
hydrogen ? 

3. What volume of hydrogen at 14° and 735 mm. will be evolved on 
dissolving 40 grams of zinc, (i) in sulphuric acid, (2) in hydrochloric acid ? 
If the same weight of iron is used instead of zinc, what volume of gas will 
be evolved ? 

4. What weight of water will be produced on burning all the hydrogen 
given off in (3) ? 

5. What volume of oxygen at 16° and 750 mm. will be required to burn 
all the hydrogen in (3) ? 

6. Describe several experiments which illustrate the principal properties 
of hydrogen. 

7. What is the action of steam on red-hot iron ? 

8. What weight of iron is contained in 60 grams of ferrosoferric oxide ? 

9. In making hydrogen I obtained 60 grams of zinc sulphate, what 
weight of zinc was dissolved ? 

10. Why is a burning taper extinguished when plunged into a jar of 
hydrogen ? 

11. How would you distinguish between hydrogen and carbon-monoxide^ 
a gas which, like hydrogen, burns with a pale blue flame in air ? 



LESSON X 

OXYGEN AND THE OXIDES 
HYDROXIDES, ACIDS, BASES, AND SALTS 

Oxygen 

Symbol O. Atomic Weight 15.9. Density 15.9 

Priestley, in England, discovered oxygen on the ist of 
August 1774, when he heated oxide of mercury by means 
of the sun^s rays concentrated by a burning glass. It was 
afterwards discovered independently by Scheele, in Sweden 
in 1775. 

Occurrence. — Of all the elements which occur on our 
planet, oxygen, either free or combined, is the most widely 
diffused and is found in the largest quantity. Oxygen occurs 
in the free state in the atmosphere, of which it forms one- 
fifth by volume. And in combination with hydrogen it forms 
f of the total weight of water on the earth's surface. It 
occurs very plentifully in a state of combination in all rocks, 
and is an essential constituent in all animal and vegetable 
structures. 

Test for Oxygen. — Oxygen gas may be recognised by 
the fact that if a splinter of wood, which has only a glowing 
spark on it, be plunged into the gas, it immediately bursts 
into flame owing to the rapid combustion. (Nitrous oxide, 
N^O, also answers this test, but we shall learn to distinguish 
it on p. 158.) 

Preparation. — (i) Oxygen is generally obtained in the 
laboratory by decomposing potassium chlorate, KCIO^, which, 

74 



LESSON X 



PREPARATION OF OXYGEN 



75 



at a temperature of about 350^, decomposes into potassium chlo- 
ride and oxygen. 

ExPT. 33. — The apparatus used is shown in Fig. 40. The 
white crystalline salt is strongly heated in a flask by means of 
a Bunsen burner. It first melts and then begins to give off 
bubbles of gas, which soon come off briskly. The degree of 
heat must be carefully regulated, or the decomposition be- 
comes so violent that the collection of the gas is difficult, and 




Fig. 40. 

sometimes the flask is burst. The following equation represents 
the reaction : — 

2 KClOo = 2 KG + 3 O2. 

For this reaction a high temperature is required, and if the flask 
is not of hard glass * it often melts before the decomposition is 
finished. 

That the chloride remaining is a different substance from the 
chlorate can be demonstrated by dissolving a little of each in 
distilled water and adding a few drops of silver nitrate solution, 
AgNOg. The chlorate is unaltered, but the chloride gives a 



* Hard glass contains potash in place of soda and is much less easily 
melted, hence it is used in cases where it has to withstand a very high tem- 
perature. 



76 



DECOMPOSITION OF HgO 



white curdy precipitate of silver chloride, AgCl, which is insolu- 
ble in nitric acid. This is a test for a chloride. 

KCl + AgNO. = AgCl + KNO3. 

ExPT. 34. — (2) In order to obtain the gas at a lower temper- 
ature, a small quantity of manganese dioxide is mixed with the 
powdered chlorate; the gas then comes off at 200^ C. before 
the salt fuses. The manganese dioxide is not altered during the 
chemical reaction and may be recovered unchanged. The part 
it plays is not thoroughly understood. 




Fig. 41. 

(3) Oxygen may be obtained by strongly heating mercuric 
oxide, wdiich decomposes into metallic mercury and oxygen. 

2 HgO = 2 Hg + O,. 

The oxide is strongly heated in the retort (Fig. 41), when the 
liquid metal collects in the receiver, whilst the oxygen is col- 
lected over water in the pneumatic trough (see also second part 
of Experiment 10, p. 9). 

(4) Manganese dioxide loses a portion of its oxygen when 
heated strongly, thus : — 

3 MnO^ = Mn.,0^ + Oo. 



LESSON X PURE OXYGEN FROM AIR 77 

The pure dioxide is placed in an iron bottle and heated to 
bright redness, when | of its oxygen escapes in the gaseous 
state. 

(5) Manganese dioxide also gives off oxygen when heated 
in a glass flask with strong sulphuric acid, 

2 MnO^ + 2 H,SO^ = 2 MnSO^ + 2 H^O + O2. 

(6) Preparation of Oxygen from the Air.* — Oxygen is 
now used in large quantities for the production of the oxy- 
hydrogen light or lime-light, and is prepared in a very interesting 
manner from the air. The method depends upon the fact that 
baryta, BaO, or oxide of barium, takes up oxygen from the 
air at a dull red-heat, but the dioxide, BaO^, thus formed is 
decomposed again at a bright red-heat into baryta and free 
oxygen, thus : — 

1st reaction 2 BaO + Og = 2 BaOg, 

2nd reaction 2 BaO^ = 2 BaO + O2. 

The baryta can be used over and over again. Instead, how- 
ever, of using two temperatures to bring about the desired 
reactions, it is found that the combination of the baryta and 
oxygen takes place when air is pumped over the heated 
baryta under pressure, nitrogen passing off; whilst, if the retort 
containing the heated baryta be afterwards exhausted by an 
air-pump, the dioxide decomposes under the lower pressure, 
the temperature remaining the same. In this way the oxygen 
absorbed from the air is pumped out of the retort, and 
after storing in a large gasholder, is forced under very great 
pressure into strong steel cylinders for transport. The amount 
of oxygen which the cylinders contain is proportional to the 
pressure under which the gas is forced into them. Owing to 
their small bulk, they are very convenient for the transportation 
of oxygen from place to place, and the gas can be used 
gradually, as required, by turning a small stopcock. In order 
to prevent the gas rushing out with great force on opening 

* Oxygen is now separated from the atmosphere by this method on a 
large scale by Erin's Oxygen Company of London ; they also separate and 
compress the residual nitrogen. 



78 PROPERTIES OF OXYGEN part ii 

the stopcock, a regulator is employed by the aid of which a 
gentle stream of oxygen can be obtained. 

(7) Oxygen may also be obtained from many other sub- 
stances containing large quantities of that element, such as 
lead dioxide, PbOg, chromium trioxide, CrOg, potassium man- 
ganate, KgMnO^, bleaching powder (which contains calcium 
hypochlorite, CaCl202), sulphuric acid, HoSO^, etc.* 

Properties. — Oxygen is a colourless, invisible, tasteless, 
and inodorous gas. It is slightly heavier than atmospheric 
air, having a specific gravity of 1. 10493 (air =1) or 15.9 
when hydrogen is taken as the unit. A litre of oxygen at 
NTP, therefore, weighs 0.0899 x 15.9 grams or i. 42941 grams. 
Oxygen is the great supporter of combustion, and bodies 
which burn in air (which contains only one-fifth its volume of 
oxygen, the rest being inert nitrogen) burn much more bril- 
liantly in oxygen. 

Combustions in Oxygen 

ExPT. 35. Phosphorus. — We may define a combustion 
as an act of chemical combination accompanied by the 
evolution of light and heat. Several gas jars of oxygen hav- 
ing been collected, a small piece of phosphorus,t about the 
size of a pea, is carefully placed in a cold brass deflagrating 
spoon which is held a little distance over a small Bunsen 
flame until it melts and takes fire. It is then at once plunged 
into a jar of oxygen, when a most vivid combustion of the 
phosphorus begins, an intensely white light being emitted 
and dense fumes of solid phosphorus pentoxide being produced, 
these, however, soon dissolve in water and are absorbed if 
a layer of water be left in the gas jar on collecting the gas 
(see Expt. 9, p. 8). If a little blue litmus solution is now 
added to the water in the gas jar after the fumes have subsided, 
it is at once turned red, showing that an acid (phosphoric acid, 
H3PO4) has been produced. 

P4 + 5 O2 = 2 P2O5. P2O5 + 3 H2O = 2 H3PO4. 

Expt. 36. — Sulphur in the same way burns with a 
* See under sulphuric acid (p. 212). f See footnote, p. 8. 



LESSON X THE OXY-HYDROGEN LIGHT 79 

bright blue flame, brighter than in air, producing a gas 
called sulphur dioxide, S0^„ which dissolves in water, as in 
the previous experiment, forming sulphurous acid, HgSOg. 

S + 02 = S0^. S02H-H,0 = H2S03. 

The presence of an acid can be shown by litmus as before.* 

ExPT. 37. — Carbon or charcoal, which burns very slowly 
and with a dull flame in air, burns brightly and with brilliant 
scintillations in oxygen,f forming carbon dioxide, CO2. If 
a little Hme-water be introduced into the gas jar after the 
combustion of carbon, it is at once turned milky owing to 
the formation of a white insoluble substance (calcium carbonate, 
CaCOg). 

C + 02 = C02. C02 + CA(OH)2 = CaCOg + H.O. 

ExPT. 38. Iron. — A bundle of fine iron wire or a steel 
watch spring can easily be burnt in oxygen, if tipped with 
burning sulphur and plunged into the gas. The oxide of 
iron formed (FegO^) drops down in the molten state and 
cracks the gas jar, unless a little sand or a layer of water 
is placed in the bottom to prevent it. 

ExPT. 39. Hydrogen. The Oxy-Hydrogen light or 
Lime-light. — When hydrogen (or, more generally, coal gas) 
is burnt in a stream of oxygen in a specially constructed 
burner (Fig. 42), an intensely hot but very feebly luminous 
flame is produced. If, however, the flame be allowed to 
impinge upon a cylinder of lime, that substance becomes heated 
to such a high temperature that it becomes incandescent and 
gives off a brilliant and dazzling white light. This is made 
use of in the magic lantern, and also in public entertainments 
where a brilliant illumination is desired. 

" Supporter of Combustion " and '^ combustible Body," 
only relative Terms. — It has been said that oxygen is a 
supporter of comb7istion, and that hydrogen is a combustible 

* Lavoisier gave the name oxygen to this gas (from o^u?^ acid, 
yevvdu)^ I produce) because of this production of acid, 
t A piece of bark charcoal shows the scintillations best. 



8o OXYGEN A "SUPPORTER OF COMBUSTION" PART II 

body^ but a little consideration will show us that these terms 
are only relative, e.g. if our rooins could be filled with hydrogen 
(or coal gas) and our gas pipes could be filled with oxygen 
(or air), we could turn on the oxygen at a gas jet, and by 
applying a light obtain a flame of oxygen burning in hydrogen. 
This would seem to reverse the ordinary state of things, for 
the oxygen would then be the combustible body, and the 
hydrogen the supporter of combustion. As a matter of fact, 
however, neither of the gases has either name rightly applied 
to it, because both are necessary to the chemical action which 
results in flame. It is by the chemical combination of the two 
substances that the heat and light are produced, and neither 
is more essential than the other for the production of the effect. 
Yet, as in everyday life, we speak of a substance that will burn 



Fig. 42. 

in air as a combustible body, it is convenient to call oxygen a 
supporter of combustion. 

ExPT. 40. — A simple experiment will illustrate this subject 
and make it quite clear. We will fill an inverted gas jar with 
hydrogen, 'and light the gas at the mouth of the cylinder, we 
will now immediately plunge a glass tube into the gas, from 
which is issuing a slow stream of oxygen from a Brin's oxygen 
cylinder furnished with a regulator, and we shall find that as 
the jet of oxygen is thrust up through the burning hydrogen, 
it takes fire and burns. We saw in Lesson IX. under hydrogen, 
that a lighted taper thrust up into that gas was extinguished, 
because the materials of the taper (carbon and hydrogen) had 
no power of combining directly with hydrogen, but oxygen has 



LESSON X THE OXIDES AND ACIDS 8i 

this power, and therefore it burns, or, more correctly, the two 
gases burn together. 

This experiment may be made more simply with coal gas 
instead of hydrogen. Common air may also be substituted 
for oxygen by delivering it slowly from a blow-pipe bellows. 



The Oxides 

All the elements, with the single exception of fluorine, are 
found to unite with or form compounds with oxygen, which 
are termed oxides. These may generally be placed in one of 
the following three classes, (i) basic oxides, (2) acid-forming 
oxides, (3) peroxides. 

Hydroxides. — Most of the oxides combine with water to 
form hydroxides, or hydrated oxides, e.g. — 

(i) Basic Oxides. Basic Hydroxides (called Bases). 

Sodium oxide Na20 + H2O = 2 NaOH Caustic soda. 

Potassium oxide K2O + H2O = 2 KOH Caustic potash. 

Calcium oxide CaO + H2O = Ca(0H)2 (Lime-water and 

slacked lime). 
Barium oxide BaO + H2O = Ba (OH) 2 Baryta water. 

Ferric oxide Fe203 + 3 H2O ~ Fe2 (OH) g Ferric hydroxide. 

(2) Acid forming Oxides. Hydroxides (called Acids). 

Sulphur dioxide SO2 -f- H2O = H2SO3 Sulphur^/^j- acid. 

Sulphur trioxide SO3 + H2O = H2SO4 Sulphur/V acid. 

Phosphorus trioxide P2O3 + 3 H2O = 2 H3PO3 Phosphor^?/j acid. 

Phosphorus pentoxide P2O5 + 3 H2O = 2 H3PO4 Phosphor/V acid. 

Carbon dioxide CO2 + H2O = H2CO3 Carbon/V acid. 

Nitrogen trioxide N2O3 + H2O = 2 HNO2 Nitr^'/^j acid. 

Nitrogen pentoxide N2O5 + H2O = 2 HNO3 Nitr/^ acid. 

Basic Oxides (or Hydroxides), when they are sokible 
in water, turn red Htmus blue, feel soapy to the skin, and all 
of them neutralise acids to form salts. 

Acids turn blue litmus red, have a sharp sour taste, and 
neutralise basic oxides and hydroxides, forming salts. 

Salts. — Basic and acid-forming oxides combine together to 
form an important class of bodies called salts. Salts may be 
formed either by the simple combination of an acid-forming 



82 ACIDS AND SALTS part ii 

and a basic oxide, or by the combination of the corresponding 
hydroxides, in which case water is eliminated, e.g.- — 

Na20 + SO2 = Na2S03 Sodium sulphzV^. 

Na20 + SO3 = Na2S04 Sodium s\i\^hate. 

BaO + SO2 = BaSOs Barium sulph//^. 

CaO + CO2 = CaCOs Calcium carbon^/^. 
2NaOH + H2SO4 = Na.,S04 + 2 HoO 
Ca(OH)2 + H9CO3 = CaCOg + 2 H9O 
Ba(OH)2 + H2SO3 = BaSOg + 2 H2O. 

Salts may be considered as acids in which the hydrogen 
has been replaced wholly or in part by a metal, and we notice 
that metals, such as Na and K, replace hydrogen atom for 
atom, these are called vionad elements ; but in the case of 
such metals as Ca, Ba, Zn, one atom of the metal replaces two 
atoms of hydrogen, these are called dyad elements. Salts may 
be produced by bringing the acid and metal into contact, e.g. — 

Zn + H2SO4 = ZnS04 + H2 
Fe + 2HC1 =FeCl2 + H2. 



Nomenclature of Acids and Salts 

It will be noticed that both the oxides of sulphur form acids, 
and the names of the two acids and their salts are examples of 
the general rule that the name of an acid produced from a 
lower oxide, ends in '^ ous''"' and its salts in '^ite,^'' whilst the 
higher oxide gives rise to an acid, the name of which ends in 
'^ ic " and its salts end in " ate^'' e.g. — 



N2O3 forms Nitr(?2/j- 


acid and ^\\xites. 


N2O5 " Nitr/^ 


" Nitrates. 


SO2 " Sulphur^z^j 


" Sulphites. 


SO3 '• Sulphur/V 


" Sulphates. 


P2O3 " Phosphor^?^j 


" Phosphites. 


P2O5 " PhosphorzV 


" Phosphates 



Acids, the names of which begin with '^ hydro, ^^ and end 
in "/^," form salts, the names of w^hich terminate in ^' ide,^'' 
thus we have 

HCl, Hydrochloric acid forms Chlorides. 

HBr, I/vdrohromic " Bromides. 

HF, Nydp-ofiuoric " Fluorides. 

H2S, IIyd?-osulphurzc " Sulphides. 



NEUTRALISATION 83 



Neutralisation. — Salts usually do not either turn red litmus 
blue, or blue litmus red, i.e. they are neutral^ but important ex- 
ceptions will be met with hereafter. 

ExPT. 41. — If we take a solution of caustic soda and add 
a solution of red litmus to it, the colour is immediately changed 
to blue. If now we add a small quantity of acid (say sulphuric 
acid), insufficient to combine wdth all the caustic soda, no 
change of colour is observed. If we then continue the addition 
of acid we come at last to a point at which the colour is neither 
blue nor red. The solution is then said to be neutral, and the 
base is said to be neutralised by an acid. At this point there 
is neither free caustic soda nor free sulphuric acid, but a 
solution of sodium sulphate. Any further addition of acid will 
now change the litmus to red. Litmus is called an indicator., 
because it indicates whether a substance is acid, neutral, or 
alkaline. 

Peroxides. — These usually contain more oxygen than the 
other two classes, a portion of which is given off on heating, e.g. 
the following ^/-oxides — 

Ba02 = BaO + O 

3 MnOo = MngO^ + O2 

Pb02 = PbO +. O. 

On treating peroxides wdth sulphuric acid, they usually decom- 
pose, a portion of the oxygen being liberated, and the residual 
basic oxide combines with the acid to form a salt, w^ater also 
being eliminated, e.g. — 

2Mn02 + 2H0SO4 = 2MnS04 + 2H..O + O2 
2Cr03 + 3H2SO4 - Cr2(S04)3+ 3H2O + 3O. 

If hydrochloric acid is used, however, chlorine may be 
liberated, e.g. — 

Mn02 + 4HCI = .MnCU + 2H2O + CI2. 

This classification is not a perfect one because some oxides 
can act either as acid-forming or basic oxides, according to the 
body with which they combine, but these must be studied under 
the head of metals. Other oxides have neither acid nor basic 
properties : e.g. H2O, N2O, NO, etc. 



84 SUMMARY AND EXERCISES 



What we have learnt 

In our tenth Lesson we have learnt the various methods of preparing 
oxygen (and particularly on the large scale from the air by Brin's process), 
its principal properties and uses and the method adopted for its storage and 
transport. The true meaning of the term " supporter of combustion." 
Oxides and hydroxides, basic and acid-forming oxides, peroxides, bases, 
acids, and salts, neutralisation and the use of litmus, and the nomenclature 
of acids and salts have been considered. 



Exercises on Lesson X 

1. Describe two methods of preparing oxygen, give equations and 
sketches of apparatus. 

2. How is oxygen prepared from the air ? How is it stored and used ? 

3. Describe experiments to prove that oxygen is a supporter of combus- 
tion, show that this term is only relative. 

4. What weight of oxygen can be obtained by heating 50 grams of 
Mn02. What will be its volume at 13^ and 720 mm. ? 

5. Give examples of basic and acid-forming oxides and hydroxides. 
What happens when these bodies are brought together ? 

6. What is meant by neutralisation ? give two examples. 

7. Give the names and formulae of the acids and the potassium salts 
derived from SO.2, SO3, N.2O3, N.2O5, P2O3, P2O5. 

8. How would you distinguish between a chloride and a chlorate ? 

9. How did Priestley obtain oxygen gas ? 

10. What compounds are formed when (i) basic oxides, (2) acid-forming 
oxides are dissolved in water ? 



LESSON XI 

Ozone 

PREPARATION — PROPERTIES AND COMPOSITION 
Formula O3. Molecular Weight 47.7. Density 23.85 

Any one who has worked an electrical machine knows that 
a peculiar smell accompanies the electrical discharge from a 
pointed conductor. This smell is due to the formation, in the 
air, of a body called Ozone, hence its name from o^etv to smelL 

Test for Ozone. — Ozone possesses remarkably energetic 
properties, hence it is sometimes called ^' active oxygen.^' 
This is shown by holding near the point of discharge, a piece 
of filter paper moistened with a colourless solution of boiled 
starch to which a few crystals of potassium iodide (Kl) 
have been added, when we observe that a deep blue colour is 
produced. This is due to the oxidising action of the ozone, by 
which iodine is liberated, and the free iodine forms a deep blue 
compound with the boiled starch. This chemical action may 
be represented as follows : — 

O3 + 2 KI + H2O - O, + I2 -f 2 KOH. 

This reaction is used as a test for ozone. 

ExPT. 42. Ozone contains Nothing but Oxygen and 
is Decomposed by Heat. — Perhaps the above-noted effects 
may be caused by some other gas besides oxygen pres- 
ent in the air. Let us, therefore, make the experiments with 
pure and dry oxygen and see whether we get the same re- 
sults. Here is a glass tube (Fig. 43) through which pure dry 

85 



86 



PREPARATION OF OZONE 



oxygen is passing, and so arranged that the gas can be sub- 
jected to the influence of an electric discharge. As the gas 
issues from the tube, whilst under the electrical influence, we 
notice the same smell as before, and our starch and iodide 
paper is at once turned blue. So we find that ozone can be 
obtained from pure dry oxygen, and it therefore contains no 
other element but oxygen. More than this, if we attach a 

piece of glass tube to the end of 
that from which ozone is issuing, 
and heat it with a Bunsen burner, 
you will see that the starch paper 
is no longer turned blue nor can any 
smell of ozone be perceived. Evi- 
dently then, ozone is destroyed by 
heat. 

(i) Preparation. — Ozone is 
best prepared by passing pure dry 
oxygen through the annular space 
between two glass tubes (one inside 
the other), the inner tube contains 
dilute sulphuric acid, in which hangs 
a platinum wire connected with 
one pole of an induction coil in 
action, whilst the outer tube is 
placed in a beaker also containing 
dilute sulphuric acid, which is con- 
nected with a wire from the other pole of the coil, as shown 
in Fig. 43. 

(2) Ozone is also produced in the electrolytic decom- 
position of water especially if the electrodes consist of fine 
platinum wire. If the oxygen given off is allowed to impinge 
upon " starch paper,'' the blue colour is at once produced. 
If electrolytic oxygen is perfectly dried, it still retains its 
smell of ozone, and on heating, it decomposes without the 
formation of water showing that the compound contains no 
hydrogen (it was at one time supposed that this compound 
was H.^Oo but the above experiment disproves it). 

(3) Ozone is also produced wiien phosphorus is allowed to 
hang, at a suitable temperature, in a flask containing moist 




LESSON XI PROPERTIES OF OZONE 87 

Properties. — Ozone is a gas possessing a peculiarly 
characteristic smell (resembling very dilute chlorine), whence its 
name. It can be condensed to a blue liquid, and is a very 
powerful oxidising agent. In its oxidising actions its volume 
undergoes no alteration, as a molecule of ozone produces a 
molecule of free oxygen, the other atom being used for the 
oxidation. Ozone produces a characteristic effect on mercury, 
taking away its lustre and mobility, destroying the convexity 
of its surface, and making it adhere to glass as a thin mirror. 
Its action on potassium iodide has already been mentioned. 
It oxidises phosphorus to phosphoric acid, HgPO^ in presence 
of moisture, sulphides to sulphates, and it has a strong oxidising 
action on all organic matter. 

Composition of Ozone. — It is now known that ozone 
is a peculiar modification or '^allotropic condition" of 
oxygen, due to a rearrangement of the atoms in the molecules, 
such that 3 molecules of oxygen (3O2), each containing two 
atoms, are rearranged so as to produce 2 molecules of ozone 
(2 O3) each containing three atoms, thus : — 

3 0, - 2 O3. 

Oxygen. Ozone. 

Experimental determination of the Composition of 
Ozone. — For many years the exact composition of ozone 
remained unknown. It was at one time supposed to be a 
compound of hydrogen and oxygen, H.^Oo, but its formation from 
pure dry oxygen as well as the fact that no water is produced 
on decomposing ozone by heat show that this is incorrect. 

(i) Ozone is Condensed Oxygen. — When a silent 
electric discharge is passed through pure dry oxygen contained 
in a glass tube furnished with two platinum wires, fused into the 
glass and ending in a bent capillary tube in which some strong 
sulphuric acid is placed (Fig. 44), it is seen that a gradual 
diminution of volume occurs, the column of sulphuric acid 
(upon which the ozone does not act) being drawn up towards 
the bulb tube. 

This diminution of bulk has never been seen to be greater 
than y\ of the volume of the oxygen (only a small proportion 
of the oxygen being converted into ozone) . 

(2) Ozone Expands -when Decomposed by Heat. — 



EXPERIMENTAL DETERMINATION OF 



II 



\ 



Fig. 44. 



Having now converted as much as is possible of the oxygen into 
ozone, let us fuse up the end of the capillary (J tube, and heat 
the bulb containing the ozone to a temperature of 300^" C, and 
after allowing the vessel to cool to the original 
temperature, we will open the sealed end of the 
U tube and we find that the volume is exactly 
what it was before the ozonisation, i.e.^ during the 
formation of ozone, oxygen is condensed, and dur- 
ing its decomposition by heat, it regains its original 
volume. Moreover, as we have seen (Expt. 42), 
the ozone is entirely decomposed by heat. 

(3) Ozone is Absorbed by Turpentine 
without Decomposition. — Certain oils, such 
as oil of turpentine, have the power of absorbing 
ozone without decomposing it, and it is found that 
the volu7ne of ozone absorbed by the oil is exactly 
double the diminution in vohane noticed when 
the oxyge7t was ozonised. It is also exactly double 
the increase of volu7ne noticed when the ozone is 
decojHposed by heat. 

Hence supposing that we ozonise oxygen and 
have a contraction of, say, i volume, and if we now absorb the 
ozone by turpentine we shall have a further contraction of 2 
volumes or a total contraction of 3 volumes. Hence 3 volumes 
of oxygen (total contraction) are condensed to form 2 volumes 
(absorbed by turpentine) of ozone. 



3 O, = 2 O3, 



therefore the formula of ozone is O3. 

(4) Ozone Destroyed by Potassium Iodide suffers 
no change in Volume. ~ When a sealed bulb containing a 
solution of potassium iodide is introduced into the ozone tube 
(Fig. 44) before ozone is produced, it is found after con- 
verting the oxygen into ozone until no further diminution of 
volume occurs that if the bulb be then broken and iodine 
liberated, no further change in the vohone of the gas is noticed ; 
whilst if the amount of the iodine liberated be exactly 
determined, and the volume of oxygen corresponding to it be 
calculated, it is found that the volume of oxygen absorbed 



LESSON XI THE COMPOSITION OF OZONE 89 

by potassium iodide is exactly equal to the diminution in bulk on 
ozonisation. 

All the above experiments are explained by the formula O3 
for ozone. 

(1) Oxygen contracts when ozonised. 

(6 vols.) 3 O2 = 2 O3 (4 vols.) decrease 2 vols. 

(2) The expansion of ozonised oxygen on decomposing it by 
heat is equal to the contractiofi on ozonising 

(4 vols.) 2 O3 = 3 Oo (6 vols.) increase 2 vols. 

(3) The diminution on ozonising and the increase on 
decomposing by heat is exactly half the volume absorbed by 
turpentine. 

Decrease and increase (above equations i and 2) = 2 vols. 
Volume absorbed by oil is 2O3 = 4 vols. 

(4) The volume of ozonised oxygen is unchanged when the 
ozone is decomposed by potassium iodide 

O3 + 2 KI + H2O = O2 + I2 + 2 KOH. 

2 vols. 2 vols. 

(before), (after). 

(5) Confirmation of Formula by the Density Deter- 
mined by Rate of Diffusion. — Another very ingenious 
experiment has proved that the formula O^ is correct. We 
have learnt on p. 63 that the relative rates of difTusion of tw^o 
gases are inversely proportional to the square roots of their 
densities. Evidently, then, if we can find the rate of dif- 
fusion of ozone compared with the rate of difTusion of chlorine 
(of known density), we can easily calculate the density of 
ozone. This has been done, and it is found that the density 
of ozone is 23.85 or ij times that of oxygen, and therefore the 
formula is O3. 

What we have leariNT 

In our eleventh Lesson we have learnt the various methods of preparing 
ozone and some of its most characteristic properties. We have seen that 
it is composed of nothing but oxygen, which is in a condensed state, such 
that three molecules of ordinary oxygen are condensed to 2 volumes of 



90 SUMMARY AND EXERCISES 



" active oxygen " or ozone which is only an allotropic modification of 
oxygen. The composition of ozone was determined by various experimen- 
tal methods and found to be O3, which formula is confirmed by the rate of 
diffusion of the gas. 

Exercises on Lesson XI 

1. How may ozone be best prepared? 

2. What are the most characteristic properties of ozone? How would 
you test for ozone ? 

3. Discuss the reasons for giving the formula O3 to ozone. 

4. I have 670 c.c. of oxygen containing 4 per cent of ozone at NTP, what 
will be the volume of the gas after heating to 300°, and cooling to 15.5° 
and 780 mm.? 

5. What weight of iodine will be liberated from potassium iodide by 
250 c.c, of oxygen containing, at NTP, 2 per cent of ozone ? 

6. A litre of oxygen at NTP, containing an unknown volume of ozone, 
liberated .0468 grams of iodine, what is the volume of the ozone and what 
percentage of ozone does the oxygen contain ? 

7. How may the composition of ozone be confirmed by a determination 
of its rate of diffusion ? 

8. Give an equation showing the action of ozone upon potassium iodide. 



LESSON XI 1 



Compounds of Hydrogen and Oxygen 
HgO and HgOg 

WATER (HVDROGEx\ MONOXIDE) - DETERMINATION OF 
ITS CHEMICAL COMPOSITION BY EUDIOMETRIC 
AND GRAVIMETRIC SYNTHESIS, AND BY ELEC- 
TROLYSIS-HYDROGEN DIOXIDE 



Water 

Formula HoO. Molecular Weight 17.9. Vapour Density 8.95 

Pure water is a clear tasteless liquid, colourless when seen in 
moderate quantity, but when viewed in bulk, possessing a bluish 
green colour. This blue colour may be observed by looking at 
a white plate, which has been lowered into a deep clear lake, or 
by looking at it through a column of distilled water about 6 to 8 
metres in length, contained in a tube with blackened sides and 
plate-glass ends. 

Chemical Composition of Water 

Cavendishes Experiments. — In the introductory lessons 
we found that we could obtain the two gases hydrogen and 
oxygen, from water, hydrogen by means of sodium and both 
gases by electrolysis. It was also seen that w^hen hydrogen 
burns in air, water is produced. 

Up to the close of last century, water was considered to be 

91 



92 



CAVENDISH'S EXPERIMENT 



an element. It was in 1781 that Henry Cavendish first made 
an exact synthesis of water and ascertained its .chemical compo- 
sition. He made very careful experiments to determine the 
proportions in which oxygen and hydrogen combine by volmne. 
An apparatus similar to that used by him is shown in Fig. 45. 
It consists of a very strong pear-shaped glass globe A, fitted with 
two wires B, for the purpose of passing an electric spark through 
the gases contained in it. 

It is fitted also with a stopcock, and it can be exhausted of 

air by the air-pump and 
screwed on the top of the 
glass gas cylinder. Cav- 
endish introduced into the 
gas cylinder two parts of 
hydrogen and one part of 
oxygen by volume. On 
opening the stopcock the 
mixed gases enter the ex- 
hausted globe, and after 
closing the stopcock they 
are fired by a spark. He 
found that after the spark 
was passed, all the gas dis- 
appeared (or '^ lost its elas- 
ticity ^') and the globe again 
became vacuous w^hilst the inside was coated with a film of dew, 
which was simply water vapour condensed to the liquid state. 

After the first charge had been fired, the globe could be again 
filled with gas by opening the stopcock, without a fresh exhaus- 
tion, and the experiment repeated until all the gas in the cylinder 
had entered the globe. In this manner Cavendish proved that 
two volumes of hydrogen and one volume of oxygen unite to- 
gether to form water, which was found to have neither taste nor 
smell, and left no sensible sediment when it was evaporated to 
dryness, — in short it w^as pure w^ater. 

Eudiometric Synthesis of "Water. — The principle which 
Cavendish used is still employed, but the modern methods 
of carrying out the experiments are much more accurate. 
The pear-shaped globe is replaced by a strong glass tube 
of uniform bore, closed at one end and called a eitdiometer 




LESSON XII EUDIOMETRIC SYNTHESIS OF WATER 



95 



(A) Fig. 46. Near the closed end two platinum wires are 
fused through the glass for the purpose of passing a spark. The 
tube is graduated in millimetres throughout its length for the 
purpose of reading off the exact height of the mercury above 
the level of that in the trough which is furnished with plate- 
glass sides. The eudiometer is first filled completely with 
mercury and inverted in the glass trough, pure oxygen 
is then introduced, and the volume, temperature, and pressure 
accurately read off. Then an excess of hydrogen is added, 
and the data read off as 



before. Next the eudio- 
meter is pressed firmly 
down on a pad of india- 
rubber fixed at the bottom 
of the trough and under 
the mercury, and a spark 
is passed. On releasing 
the eudiometer the volume 
of gas is found to have 
diminished, and all the 
oxygen has disappeared 
together with exactly twice 
its volume of hydrogen. 
For the purpose of ascer- 
taining the temperature 
of the gases, a thermo- 
meter is hunor near the 




Fig. 46. 



eudiometer, and an interval of rest is allowed between each 
reading so as to allow the whole apparatus and the air of 
the room to reach an equable temperature. The pressure to 
which each volume of gas is subjected is found by reading 
the height of the barometer and subtracting the height of 
the column of mercury in the eudiometer above the level of 
that in the trough ; for it is evident that when at rest the 
pressure inside the eudiometer at the level of the mercury 
in the trough is equal to the atmospheric pressure shown by 
the barometer, but this is made up of the pressure of the gas 
and the pressure of mercury in the eudiometer, and therefore 
by subtracting the latter, which can be easily read off, we 
obtain the pressure of the gas. The volumes of the gases at 



94 



COMPOSITION OF STEAM 



PART II 



each reading being known they are reduced to o° and 760 
mm. NTP (normal temperature and pressure) before begin- 
ning further calculations. The volume of liquid water pro- 
duced is almost inappreciable, a correction can, however, be 
made for it, and also for the tension of aqueous vapour 
(see p. 105). 

To render this clear let us suppose that the amount of 
oxygen was 50 vols, at NTP, and that of hydrogen, 150 vols., 
and that after passing the spark a residual 50 vols, at NTP, con- 
sisting of pure hydrogen was found, then it is evident that all 
the oxygen 50 vols, has combined with 150 — 50=100 vols, of 
hydrogen ; or the ratio of oxygen to hydrogen by volume is 1:2. 
Volumetric Composition of Steam by Synthesis. — 
Gay-Lussac was the first to ascertain that two volumes of 

hydrogen and one volume of oxy- 
gen combine together to form two 
volumes of gaseous steam. This 
can be proved by means of the 
apparatus shown in Fig. 47. It 
consists of a eudiometer tube, which 
is surrounded by a wider glass tube 
or jacket, through which can be 
passed a current of vapour of amyl 
alcohol (which boils at 132° C.) 
from the flask in which the liquid is 
boiled, the object of this being to 
keep, in the state of vapour, the 
whole of the w^ater w^hich is formed 
by the union of the oxygen and 
hydrogen. The vapour of the amyl 
alcohol passes out and condenses in 
^^s- 47- the flask placed in cold water. The 

eudiometer is joined to another plain tube open at the top, the 
whole forming a U tube which has a branch tube at the side 
for the inlet and outlet of mercury, thus enabling the pressures 
to be adjusted by raising or lowering the mercury reservoir. A 
mixture of exactly two volumes of hydrogen and one of oxygen 
obtained in the apparatus shown in Fig. 49, is first passed into 
the eudiometer and its temperature raised to 132° by means of 
the amyl alcohol vapour. The mercury in both limbs is adjusted 




LESSON XII 



ELECTROLYSIS OF WATER 



95 



to the same level and the volume read off, after which the pres- 
sure is reduced and the spark passed. Combination takes place, 
and it is found, on again adjusting the pressures, that the 
volume of gaseous water produced (measured at the same 
temperature, viz. 132°) is exactly f of that of the original 
mixture, showing that two volumes of hydrogen and one of 
oxygen combine to form two volumes of gaseous steam, a 
cont7'action of \ of the bulk of the 
uncombined gases taking place after 
combination. 

ExPT. 43. Volumetric Analysis 
of Water. — Fig. 48 shows the ap- 
paratus which may be used for the 
electrolytic decomposition of water. 
The bulb and the straight tube are 
filled with acidulated water, which, on 
opening the two stopcocks at the end 




Fig. 48. 

of the U tubes rises and fills the gas collecting tubes. A current 
of electricity from a battery is then passed through the liquid by 
means of the two platinum wires fused through the two tubes 
forming the (J- The platinum wires terminate in platinum plates 
or electrodes from which, on joining up the connecting wires, a 
constant stream of minute bubbles of gas rises, oxygen from 



96 



ELECTROLYTIC GAS 



PART II 



the -f and hydrogen from the — electrode, and the gases collect 
separately in the stoppered tubes. 

It will be noticed that the volume of oxygen given off is 
rather less than half that of the hydrogen ; and the reason is 
that (i) oxygen is more soluble in water than hydrogen and 
a portion remains dissolved, and (2) a little ozone is formed 
which occupies less space and is more soluble than oxygen. 
That the larger volume of gas is hydrogen may be shown by 
opening the stopcock when the pressure of liquid in the bulb- 
tube forces the gas out and it may be lighted and observed to 
burn with the characteristic flame of hydrogen. The other gas 
ignites a glowing chip and is oxygen. That the gas contains 
ozone, may be demonstrated by allowing the issuing gas to im- 
pinge against an iodised starch 
paper (see ozone, p. 85) which is 
turned blue. By raising the tem- 
perature of the acidulated w^ater to 
100° the solubility of the oxygen 
is diminished, whilst that of the 
hydrogen remains practically un- 
changed, and at the same time 
the formation of ozone is avoided, 
so that the true volume relation is 
thus more closely attained. 

Electrolytic Gas. — This is 
the name given to the mixture of 
hydrogen and oxygen in the pro- 
portion 2:1, produced in the elec- 
trolysis of water, when they are collected together. Fig. 49 
shows a form of apparatus used for making electrolytic gas, 
which explains itself. 

ExPT. 44. Experiments -with Electrolytic Gas. — When 
electrolytic gas is brought in contact with a flame or electric 
spark, combination takes place with explosive violence, owing 
to the sudden expaiisiofi of the gaseous steam produced by 
the great heat evolved in the act of chemical combination. 
This is best shown by collecting a small flask of electrolytic 
gas, and fitting it on an india-rubber stopper pierced with two 
holes, through which pass two stout copper wires terminating 
in small pieces of platinum wire not quite in contact. The 




Fig. 49. 



LESSON XII GRAVIMETRIC SYNTHESIS OF WATER 



97 



flask is supported as shown in Fig. 50, and covered with a 
wire gauze covering. When a spark from an induction coil 
is passed, a loud explosion takes place which shatters the flask 
to fine dust. 

Soap bubbles may be blow^n with electrolytic gas, and as they 
ascend they may be touched with a lighted taper giving rise to 
a loud explosion. Care must, however, be taken that the light 
does not come in contact w^ith the end of the tube which delivers 
the explosive gases. 




Fig. 50. 



Gravimetric Synthesis of "Water. — Knowing the com- 
position of water by volume, and also the densities of the two 
gases, it is easy to calculate the composition of water by weight. 
but it is important to be able to verify the result by actual 
experiment. 

The method w^hich w^as first proposed and carried out by 
Berzelius and Dulong in 1820, consists in passing pure hydro- 
gen over a weighed quantity of red-hot copper oxide, CuO, 
w^hen the oxide is redticed or robbed of its oxygen, which goes 
to form water with the hydrogen, and leaves metallic copper 
behind, thus : — 

CuO + H, = Cu + H>0. 



GRAVIMETRIC SYNTHESIS 



By weighing the copper oxide before the experiment and the 
copper remaining after the experiment, the weight of oxygen 
used is found. Then by carefully collecting and weighing the 

water produced, and subtracting the 
oxygen previously found, the weight 
of hydrogen used is obtained. The 
above experimenters found as a mean 
of four experiments that water w^as 
composed of — 




Found. 
Oxygen 88.901 
Hydrogen 11.099 



Calculated. 
88.864 

100.000 



In 1843 Dumas and Stas made 
a most careful repetition of these 
experiments, taking every conceiv- 
able precaution against experimental 
errors. The hydrogen which they 
i obtained from sulphuric acid and 
' zinc w^as passed through a series of 
U tubes shown in Fig. 51, contain- 
ing nitrate of lead, sulphate of silver, 
caustic potash, and phosphorus pent- 
oxide which serve to purify the gas 
from traces of sulphuretted hydro- 
gen, arseniuretted hydrogen, sulphur 
dioxide, oxides of nitrogen, and 
moisture. The pure hydrogen thus 
obtained was passed over the weighed 
bulb A (Fig. 51) containing copper 
oxide and joined to the bulb B 
which is also weighed and serves for 
the collection of the water produced, 
any water vapour being absorbed by 
the following weighed U tubes . They 
found by means of nineteen sepa- 
rate experiments that 840.161 grams of oxygen were consumed 
in the production of 945.439 grams of water, or the percentage 
composition of water by weight is 



lESSON XII HYDROGEN DIOXIDE 99 



Oxygen 88.864 

Hydrogen 11.136 



which agrees exactly with the composition calculated from the 
volumetric analysis. 

Lavoisier's Experiments, — At the end of last century 
Lavoisier examined the composition of water by the following 
method, depending on the fact that when water vapour is passed 
over red-hot iron the oxygen is absorbed, and the hydrogen 
passes on in the free state. 

3Fe + 4H2O =Fe30, + 4H,. 

The apparatus used by Lavoisier is of great historical interest. 
The water was allowed to drop slowly into a tube from which it 
flowed into a gun barrel, heated in a furnace. Here the decom- 
position took place, and the hydrogen passed on and was col- 
lected and measured in a large bell jar, whilst any undecomposed 
steam was condensed, and the water collected and weighed. The 
results thus obtained are influenced by so many experimental 
errors that no very exact numbers were obtained, but still the 
results were sufficiently near the truth for the state of the science 
more than one hundred years ago. 



Hydrog-en dioxide (or peroxide), H2O2. 

Preparation. — Hydrogen dioxide is an unstable compound only 
known in solution, and has not been prepared in the pure state, it is 
obtained by the action of dilute sulphuric acid on barium dioxide, 

Ba02 + H2SO4 = BaS04 + H2O2. 

Properties. — Hydrogen dioxide is so unstable a compound that it is 
easily decomposed into oxygen and water on slightly raising its tempera- 
ture, 

H2O2 - H2O + O. 

It possesses bleaching properties like chlorine, although acting more 
slowly, and is sometimes used for bleaching hair and old engravings. It 
acts as a powerful oxidising agent, converting ferrous salts into ferric salts 
in presence of an acid, and oxidising black lead sulphide, PbS, to white 
lead sulphate, PbS04. 

Hydrogen dioxide contains an atom of oxygen very loosely combined, 



HYDROGEN DIOXIDE 



and this often abstracts another atom of oxygen, which is also loosely 
combined, from many other compounds, thus giving rise to the more stable 
molecule of free oxygen. It thus apparently acts as a reducing agent, that 
is it abstracts oxygen from other compounds. Thus ozone is decomposed 
into ordinary oxygen, 

03 + H202 = 2 02+H20. 

Silver oxide is reduced to the metallic state whilst oxygen is set free. 

Ag20 + H2O2 = Ag2 + H2O + O2. 

It also 7'educes manganese dioxide and potassium permanganate in 
presence of sulphuric acid, manganous sulphate being formed and oxygen 
liberated, 

MnOs + H2SO4 + H2O2 = MnS04 + 2 H2O + O2. 
2 HMn04 + 2 H2SO4 + 5 H2O2 = 2 MnS04 + 8 H2O + 5 O2. 

The potassium permanganate, KMn04, may be supposed to be decom- 
posed by the sulphuric acid, H2SO4, and the free permanganic acid, 
HMn04, liberated, 

2 KMn04 + H2SO4 = K2SO4 + 2 HMn04. 

This is then reduced to maganous sulphate by the hydrogen peroxide 
and sulphuric acid, in accordance with the above equation. 

What we have learnt 

In our twelfth Lesson we have studied the various experimental methods 
by which the composition of water has been ascertained. We have seen 
that Cavendish made the first synthesis of water; he showed that two 
volumes of hydrogen and one volume of oxygen unite together to form 
water, and nothing else but water. The more exact eudiometric method 
is described in detail. We have learnt also that when two volumes 
of hydrogen and one volume of oxygen are combined at a temperature of 
132'^ two volumes of steam are produced. The electrolytic analysis of 
water and the explosive properties of electrolytic gas are next described. 
We have seen how the volumetric methods are confirmed by gravi- 
metric experiments, in which water is produced synthetically by passing 
pure hydrogen over heated copper oxide. The water produced is weighed, 
and the weight of oxygen it contains is found by noting the loss in weight 
of the copper oxide. The weight of hydrogen is then obtained by dif- 
ference. 

The preparation and principal properties of hydrogen dioxide have been 
described ; it is shown that this compound acts as a strong oxidising agent, 
and yet in many reactions it apparently acts as a reducing agent, owing 
to its containing an atom of oxygen loosely combined. This unites with 
another loosely combined atom of oxygen to form a stable molecule of free 
oxygen. 



LESSON XII SUMMARY AND EXERCISES 



Exercises on Lesson XII 

1. How did Cavendish prove the composition of water ? 

2. Describe in detail the method of making a eudiometric synthesis of 
water. 

3. What is the volumetric composition of steam, and how is it experi- 
mentally determined ? 

4. By what experimental method has the gravimetric composition of 
water been determined ? 

5. The density of oxygen being 15.9, calculate the gravimetric compo- 
sition of water from its volumetric composition. 

6. In an experiment on the gravimetric composition of water 0.5278 
grams of water were obtained. What weight of copper oxide had been 
decomposed, and what volume of hydrogen at 15° and 746 mm. had been 
used to decompose it ? 

7. How is hydrogen dioxide prepared ? Explain its reducing action on 
Ag20, O3, KMn04, and Mn02. 

8. What weight of oxygen is contained in 500 c.c. of water measured 
at 40? 

9. Five grams of water are decomposed by electrolysis. What volume 
of electrolytic gas measured at 15° and 745 mm. pressure has been ob- 
tained ? 

10. What volume of liquid water is produced by the explosion of 
100 c.c. of electrolytic gas (at NTP) in a eudiometer ? (i c.c. of water 
weighs a gram.) 

11. In an experiment on the gravimetric synthesis of water, 150 grams 
of zinc were dissolved in dilute acid, and the hydrogen passed over heated 
copper oxide : what weight of copper oxide would be decomposed, and 
how much water formed ? 



LESSON XIII 
HEAT RELATIONS OF WATER 

EXPANSION AND CONTRACTION — POINT OF MAXIMUM 
DENSITY — TENSION OF VAPOUR — EVAPORATION — 
MELTING AND BOILING POINTS — LATENT HEAT- 
FREEZING MACHINES — SPECIFIC HEAT — DULONG 
AND PETIT'S LAW 

Expansion and Contraction of Water. Point of Maxi- 
mum Density. — If we cool water from loo'^ to o^, it is 
found to contract until the temperature 4^ is reached (or more 
exactly 3.945), if it is cooled below this point // expands in 
passing from 4° to 0°, whilst when the water at 0° becomes ice 
at 0° (when it freezes), a sudden and great expansion takes 
place. 

It is evident, that since water contracts until it reaches 4°, 
and further cooling causes an expansion, that 4° is the point 
at which water is denser than at any other, i.e. it is the point 
of maximum density, and is heavier bulk for bulk at this tem- 
perature than at any other. 

Dr. Hope's Experiment. — This is w^ell shown by an 
experiment first made by Dr. Hope, which is here slightly 
modified. Fig. 52 shows a cylindrical vessel, perforated laterally 
by two holes, in which thermometers are fixed, one at the top 
and the other at the bottom. Round the centre of the vessel 
is a shelf containing broken ice. When the vessel is filled with 
water at the ordinary temperature of the air, say 15°, it is 
noticed that the lower thermometer falls rapidly to 4°, at which 



LESSON XIII 



MAXIMUM DENSITY 



103 



point it remains stationary, whilst the upper one is scarcely 

affected. This is because the water cools, and contracts, and 

becomes heavier, and therefore sinks to the bottom. This goes 

on until the lower half of the w^ater reaches 4°, when no further 

circulation in that portion takes 

place. The upper half is scarcely 

affected because water is a very 

bad conductor of heat, but the 

middle portion of water is soon 

cooled below 4°, and becomes 

lighter, and therefore ascends and 

sets up a circulation in the upper 

half of the water causing the upper 

thermometer to begin to fall. It soon 

indicates a temperature of 0°, whilst 

the lower one remains stationary at "^^ ^ ^^J^ ^ fcz^^^ ^^ W 




4°, showing that the water at that 



Fig. 52. 



temperature is heaviest, sinking to 

the bottom, and that the water at 0° is lighter because it rises 

to the top. 

Increase in Volume on Solidification. — When water 
at o^ freezes and becomes ice at 0°, the change of state is 
accompanied by a sudden and great expansion amounting to 
about 10 per cent: thus i vol. of water at 0° becomes 1.102 
vols, of ice at 0°, or an increase in volume of 10.2 per cent. 
The following table shows the density or specific gravity of ice 
as compared with water at 0^-25° C. (water at 4^ = i). 



Temperature, 


Density. 


Temperature. 


Density. 


Ice 0^ 


0.91674 


Water 8^ 


0.99987 


Water 0^ 


0.99987 


" 10^ 


0.99973 


" 2'' 


0.99996 


.. ,50 


0.99912 


" 4^ 


1.00000 


" 20^ 


0.99821 


u 6° 


0.99997 


" 25° 


0.99708 



Effects in Nature. — The force with which water expands 
on freezing is almost irresistible. This is evidenced by the 
bursting of strong iron water-pipes in winter. It is popularly 
believed that the t/iaiu bursts the pipes, but this is an erroneous 
conclusion from the fact that the pipes first begin to leak when 
the thaw takes place. It is evident that at the time of bursting, 



I04 



EFFECTS OF FROST 



PART II 



the pipes are filled with solid ice and none can therefore escape. 
It is only when the ice melts during the subsequent thaw, that 
any leakage of water is seen. 

If we examine the effects of the point of maximum density 
of w^ater being 4°, we see that it is owing to this fact as well 
as to the expansion on solidification that ice always forms 
on the suf'face of ponds and lakes. This is a most beneficent 

provision of nature, for 
WTre it not so, our lakes 
and rivers would, in 
winter, freeze at the 
bottom first, and w^ould 
soon become a solid 
mass of ice which even 
all the warmth of sum- 
mer would probably be 
unable to thaw. As it 
is, the whole mass of 
water in a lake must 
first reach 4° before 
the surface layers can 
freeze ; hence very deep 
lakes seldom freeze. 
The covering of ice 
protects the deeper 
water from the exces- 
sive cold, and fish and 
other aquatic life is 
preserved through the 
severest winter, the 
lowest depths never 
falling below 4^ C. 

The effects of frost 
in the disintegration of 
the soil and rocks can now be easily understood. The gravel 
path w^iich is saturated with water before a frost, is lifted bodily, 
because all the interstices filled with water are made larger when 
the water is turned into ice. Hence a thaw^ leaves the path in 
a soft and pulpy condition, because the ice melts and sinks into 
the ground, leaving the gravel slightly upheaved and very 




Fig. 53. 



LESSON XIII VAPOUR TENSION 105 

porous. Rain water, freezing in a crevice in the rocky cliffs 
of our coasts, exerts an enormous force, rending the rocks 
and making the gap wider; this action, in time, is sufficient 
to detach large masses of rock, which are thus hurled down 
below. 

Tension of Aqueous Vapour. Evaporation. — Water, 
even at the ordinary temperature of the air, is constantly giving 
oflf vapour, and we know that if a saucer full of water is left 
in the open air for a few days it gradually evaporates into 
the atmosphere. This vapour, which comes off from w^ater 
(and even from solid ice) at all temperatures, exerts a pres- 
sure if it is confined ; and this pressure is constant for each 
temperature. This is known as the elastic force or tension of 
aqueous vapour, and may easily be seen by taking two exactly 
similar barometer tubes, and into the Torricellian vacuum of 
one of them, introducing a few drops of water. The mercury 
in this one wdll be depressed, the amount of the depression 
being the greater the higher the temperature. The following 
numbers show the depression in millimetres of mercury, for 
water vapour at different temperatures : — 



Tension of Aqueous Vapour 



Temperature. 


Tension. 


Temperature. 


Tension 


Ice at + 10° C. 


2.09 


40° 


54.90 


0° 


4.60 


60^ 


148.79 


Water at + 10° 


9.16 


80° 


354.28 


20° 


17.39 


100° 


760.00 


•• 30^ 


31-55 


101° 


787.63 



The above facts explain the more rapid evaporation of water 
in summer ; another factor, however, being the dryness of the 
atmosphere. Evidently then, when water is heated, the ten- 
sion of its vapour increases, until at 100° its tension is equal to 
the atmospheric pressure ; and if water is heated to 100° in a 
Torricellian vacuum, when the atmospheric pressure is 760 mm., 
the mercury is depressed to the level of that in the trough. 
The apparatus for showing this will be readily understood from 

Fig. 53- 

The vapour tensions of other liquids are different from those 
of water at the same temperature, e.g. at 0° C, they are 



io6 



VAPOUR TENSIONS 



Vapour Tensions at O^ 



Water . 

Alcohol . . . . 
Carbon bi-sulphide 
Ether . . . . 
Sulphur dioxide (liquid) 



4.6 mm. 

13 mm. 

132 mm. 

182 mm. 

1 165 mm. (boils below o^) 



The vapour tensions of water, alcohol, and ether are clearly 
shown in Fig. 54. The first is a 
barometer tube wdth a Torricellian 
vacuum ; into the other three have 
been introduced ist, water; 2nd, 
alcohol ; 3rd, ether, so as to satu- 
rate the space with their vapour, and 
produce the maximum tension for 
that temperature. 

ExPT. 45. Melting and Boil- 
ing Points of Water. Latent 
Heat. — If we fill a beaker with 
pounded ice and apply heat by a 
Bunsen burner, we notice that the 
thermometer placed in it indicates 
0° C, and although heat is con- 
stantly being given to the ice, no 
increase of teiJiperatiire takes place 
until the ice is melted, the only effect 
of the heat is to melt the ice or 
change its state from solid to liquid. 
This heat which is absorbed with- 
out raising the temperature is called 
" the latent heat of fusion.*" When 
all the ice is melted, the tempera- 
ture of the water rises until 100^ is 
reached, at which point the water 
boils. When this point is reached 
no further increase of temperature 
takes place, any additional heat 
merely causing a change from liquid 
to gaseous water. Here again, the 

heat required to vaporise water is rendered latent, and is called 

the ''latent heat of vaporisation.*" 




Fig. 54. 



LESSON XIII 



MELTING AND BOILING POINTS 



107 



Influence of Pressure on the Melting and Boiling 
Points. — We shall see that a liquid boils when the tension 
or pressure of its vapour is equal to the pressure above it, 
hence it is only correct to say that water boils at 100° C, 
when the atmospheric pressure is at its normal, viz., 760 mm. 
Under diminished pressure water boils at a lower temperature, 
and under increased pressure w^ater boils at a higher tempera- 
ture than 100°. Hence it is important, in graduating a ther- 
mometer, to read off the exact pressure at which the boiling- 
point is marked on the stem. On Mont Blanc w^ater boils at 
85^, and eggs cannot there be cooked by boiling in an open 
vessel. Whilst under a pressure of two atmospheres water 
does not boil until 120.6° is reached, and at twenty atmo- 
spheres' pressure water boils at 213°. Knowdng the tempera- 
tures at which water boils under different pressures, it is possible 
to measure the heights of mountains by means of the thermo- 
meter. It is only necessary to determine the temperature at 
which water boils at the summit and at the base or the sea-level. 
From these determinations the atmospheric pressure can be 
deduced, and from this the 



height of the place at which 
the experiment was made. 

ExPT. 46. — A simple 
experiment will show that, 
under diminished pressure, 
water boils at a lower tem- 
perature ; and, paradoxical 
as it may seem, it is possible 
to boil hot water by cooling 
it. Fig. 55 shows the method 
of doing this. Water is 
boiled rapidly in a strong 
round-bottomed glass flask, 
furnished with a stopper, 
through which an ordinary 
gas tap is passed ; whilst 




Fig. 55- 



the water is boiling, the tap is closed, the heat at the same 
moment being withdrawn. The flask is now inverted in a 
retort-stand ring, and a shower of cold water poured over it 
from a sponge. The water immediately begins to boil afresh. 



io8 LATENT HEAT 



because the steam in the upper part of the flask is thus con- 
densed, and the pressure on the hot water released. Under 
these conditions the water is still hot enough to boil under the 
diminished pressure. 

In the same way the melting point of ice is influenced by pressure, but 
to a much smaller extent; under increased pressure, the melting point of 
ice (and of all substances which expand on solidifying) is lowe.red. Under 
a pressure of 16.8 atmospheres ice melts at — 0.126^ C. 

On the other hand, bodies which contract on solidification, such as 
sulphur, wax, etc., exhibit an opposite change, for the mehing point is 
raised. 

Definition of Boiling Point. — We may now^ define the 
boiling point of water (or of any other liquid) as that point at 
which the tejtsio7i of its vapour is equal to the superincmnbent 
atinospheric pressto'e^ and it is clear that water boils exactly at 
100° only when the atmospheric pressure is 760 mm. 

Latent Heat of Fusion. Hotv it is Measured. — When 
steam condenses to liquid water, its latent heat is given off as 
sensible heat. Similarly, when water freezes to solid ice, its 
latent heat of fusion is given off as sensible heat, and in these 
operations just as much heat is given out as is required to 
perform the opposite change of water into steam, or ice into 
water. 

ExPT. 47. Ice on melting Absorbs Heat. — If we take 
a kilogram of water at 80° and mix it with a kilo of liquid 
water at 0°, the result is 2 kilos of water of 40°, or i kilo in 
falling from 80° to 40^ has lost just as much heat as is re- 
quired to raise it from o" to 40°. But if we mix a kilo of 
water at 80° (more exactly 79.25°) with a kilo of ice at 0°, 
the result is 2 kilos of water at 0°. In other words, the hot 
water has given out just sufficient heat to vielt the ice, but 
none in addition to raise the temperature of the liquid w^ater. 
Or ice, in melting, requires as much heat as would raise its 
own weight of water through 80°, or 80 times its weight 
through 1°. 

The Thermal Unit or Calorie. — Hence the latent heat 
of fusion of water is 79.25 thermal units ; a thermal unit being 
the amount of heat required to raise a kilogram of water from 
0° to 1° C, this is called a Calorie, 



LESSON XIII THE THERMAL UNIT 109 

[In France the thermal unit is that quantity of heat which is necessary 
to raise i kllograin of water from 0° to i"^ C, and is called the Calorie. In 
England the thermal unit is that amount of heat which will raise i pound 
of water from 0° to i'^ C, and is called the pound-degree unit. 

I thermal unit (Enghsh) =0.45 Calorie.] 

ExPT. 48. Water on Freezing gives out Heat. — 
That the latent heat of fusion is given off by water, when it 
passes from the liquid state at o^ to the solid state at 0°, 
may be shown by placing a metal vessel containing cold 
water in a bath of mercury, at a temperature say of— 15^.* 
By placing a thermometer in the mercury bath, and also in 
the water, it will be found that whilst the temperature of the 
water falls to o*^, that of the mercury rises ; but whereas, during 
the passage of the liquid water into solid ice, the thermometer 
placed in it shows no alteration, but remains constant at 0°, 
that in the mercury co7itmiies to rise rapidly^ indicating that 
the latent heat is being given out by the water as sensible 
heat. 

Latent Heat of Steam. — The latent heat of steam is 
given out as sensible heat when steam condenses to water, 
and this affords a means of ascertaining its amount. 

ExPT. 49. — If we pass steam at 100° into a kilogram of 
water at o^ until the water is raised to 100°, it will be found 
that the weight of water is now 1.187 kilos, or 0.187 kilo of 
steam at 100° has given out sufficient latent heat in condensing 
to raise i kilo of ice-cold water through 100°, or i kilo of steam 
would raise 5.36 kilos of water from o^ to 100°, or 536 kilos 
through 1°. Hence the latent heat of steam is said to be 536 
thermal units or Calories. 

The apparatus for showing this Experiment will be readily 
understood from Fig. 56. 

Water Frozen by its own Evaporation. — We have 
already seen that when liquid water (or any other liquid) 
passes into the gaseous state, a large amount of heat is absorbed, 
or becomes latent, in order to keep it in that state. Not only 
is this true of water when it is boiled by the application of heat, 
but it is also true of water which slowly evaporates. 

* Temperatures below the freezing point are indicated by a minus sign 
placed before the figures. 



CARRE'S FREEZING MACHINE 



ExPT. 50. — A beautiful experiment will illustrate this. 
Fig. ^'] shows Wollastoirs Cryophorus, which .consists of two 

glass bulbs connected 




Fig. 56. 



by a glass tube. Water 
is placed in B, and this 
is briskly boiled until 
the steam has driven 
out all the air in both 
bulbs through the open 
tube at the lower end of 
A ; this is then sealed 
up by the blowpipe, 
the lamp w^hich boils 
the water being at the 
same instant withdrawn. 
After the apparatus has 
cooled, all the water is 
allowed to run into B, 
whilst A is placed in a 
freezing mixture. The 
tension of aqueous 
vapour is now much 
reduced in A, when 



aqueous vapour passes over from B to A so rapidly (tending to 
produce an equilibrium of pressure), and so much heat is 
absorbed from the water by the vapour passing over, that the 




remaining water is cooled down below o^\ and freezes to a solid 
mass of ice. 

Carre's Freezing Machine. — This principle has been 
applied on a practical scale for the production of ice. Fig. 58 
shows an apparatus which is known as Carre's Freezing 
Machine. 



LESSON XIII 



FREEZING MIXTURES 



Water in the strong bottle C is in communication with a 
reservoir B, containing strong sulphuric acid, which has the 
power of absorbing water vapour rapidly. The reservoir can 
be exhausted of air by the air-pump A. After working the 
pump for a few minutes, the pressure in C is so much reduced 
that the water begins to boil, and so much heat is rendered 
latent in the vapour which passes over into B, that the water 
in C is soon cooled below o°, and freezes to a solid mass of ice. 
The tap is now closed, and the bottle removed. 

Freezing Mixtures. — In many cases, when salts are mixed 
with water, their solution is accompanied by a diminution of 
temperature owing 
to the absorption 
of sensible heat 
which becomes 
latent during the 
passage of the salt 
from the solid to 
the liquid state, e.g. 
when 500 grams of 
potassium sulpho- 
cyanide, KSCN, is 
mixed with 400 
grams of cold 
water,the tempera- 
ture sinks to — 20°. 
When 32 parts of 
common salt, 
NaCl, is mixed 
with 100 parts of pounded ice or snow, the two solid bodies 
become liquid, and so much heat is rendered latent that the 
temperature sinks to — 23°. Equal weights of crystallised 
calcium chloride, CaCl2.2 H2O, and snow, when mixed together, 
give a freezing mixture, the temperature of which sinks from 0° 
to - 45"^ C. 

In the case, however, of mixing many anhydrous salts with 
water a chemical combination takes place, in which heat is 
evolved, e.g, anhydrous calcium chloride, CaCl,, becomes hot 
when mixed with water, because this salt has such a great 
chemical attraction for water. 




Fig. 58. 



SPECIFIC HEAT part ii 



Specific Heat. — We have seen that the thermal unit is that amount 
of heat which will raise a kilo of water through i°, and has reference to 
water only. A unit weight of iron or copper would require a differe^it 
amount of heat to raise its temperature i°. Different bodies have different 
capacities for absorbing heat, and they give out different amounts of heat 
on cooling through the same interval of temperature. This is expressed 
by saying that they have different specific heats {or calorific capacities). 
The specific heat of a substance is the ratio between the amount of heat 
required to raise its temperature through a certain interval, compared loith 
the amou7it of heat required to raise the same weight of water through the 
same interval. It is expressed by the number of thermal units required to 
raise a unit weight of the substance through i^ of temperature, e.g. only 3V 
as much heat is required to raise i gram of mercury i*^ as is required to 
raise i gram of water 1° ; hence, we say, the specific heat of mercury is 3^^ or 
i33. The specific heat of water is taken as the standard with which all 
others are compared, e.g. 

Specific Heats 

Water 1.0000 Lead 0.0314 

Copper 0.0939 Mercury 0.0332 

Bromine 0.0843 Iodine 0.0541 

Arsenic 0.0822 Sodium 0.2930 

Dulong and Petit's Law. — It was found by Dulong and 
Petit that the product of the specific heat into the atomic weight 
of solid elements is approximately constant, being about 6.3. 
Hence the specific heat of an element is sometimes of value as 
a control in determinations of atomic weight. The following 
will serve as examples : — 



Specific Heat 




Atomic 




Atomic Heat 


(of Equal Weights). 


Weight. 


(OF 


Ato 


MIC Weigh 


Antimony . 0.0513 


X 


1 19.4 


= 




6.12 


Bismuth . . 0.0308 


X 


206.4 


= 




6.35 


Cadmium . 0.0567 


X 


111.3 


= 




6.31 


Gold . . . 0.0324 


X 


195.7 


= 




6.34 


Iron . . . 0.1138 


X 


55.6 


= 




6.32 


Zinc . . . 0.0956 


X 


65.0 


= 




6.21 



In other words^ if instead of taking equal weights of these 
elements for the determination of their specific heats, we take 
atomic weights^ then we find that the specific heat (of the 
atom) is constant, or the same a7no2iiit of heat is 7ieeded to heat 
an ato77i {or a7i equal 7i7C77ibe7' of ato77is) of a7iy of the solid ele- 
77ie7its to the sa77ie exte7it. 



LESSON XIII SPECIFIC GRAVITY 113 

It is evident from the above that the specific heats of the solid elements are 
inversely proportional to their atomic weights, or 

Specific heat =-: : r-r- X 6.3 

Atomic weight 

approximately, and therefore we find that the 

6.3 



Atomic weight - 



Specific heat. 



From this last equation it is easily seen how the determination of specific 
heat may be, and has been, used as a control in the determination of the 
atomic weights. 



Specific Gravity or Relative Density 

Not only is water taken as the standard substance by which 
to compare specific heats and to define thermal units, but, just 
as hydrogen is taken as the standard with 
which to compare the densities of all gases, 
so water is taken as the standard with which 
to compare the densities or specific gravities 
of all liquid and solid bodies. When we say 
that the specific gravity or relative density 
of cast iron is 7.5, we mean that, volume 
for volume, the metal is seven and a half 
times as heavy as water, or a piece of cast- 
iron whose volume is a cubic centimetre 
weighs 7.5 grams, for we learnt on p. 39 that 
a cubic centimetre of distilled water at 4^ 
weighed i gram. When we say that the 
specific gravity of ether (Cg H.) 2O is 0.72, we mean that it is 
only about three-quarter times as heavy as water. The metal 
sodium is lighter than water, its specific gravity being 0.974, 
whilst that of pine wood is 0.5. 

Use of Specific Gravity Bottle. — When we want to find 
the specific gravity of a liquid, the simplest method is to com- 
pare the weights of exactly equal volumes of that liquid and 
water. For this purpose the specific gravity bottle is used 
(Fig. 59) ; these are generally adjusted to hold exactly 50 grams 




114 SPECIFIC GRAVITY PART ii 

of distilled water when filled at 60° F., the mean temperature 
of the air. Then all that is necessary is to weigh the clean and 
dry specific gravity bottle filled with the liquid to be examined, 
taking care that it has a temperature of 60°, and is completely 
filled, so that when the perforated stopper is inserted, the excess 
of liquid emerges from the hole in the stopper, and is then care- 
fully removed. 

Specific Gravity of Solid Bodies. — There are many 
methods of finding the specific gravity of solid bodies, but 
nearly all depend upon the principle of Archimedes, viz., that 
when a solid body is weighed whilst fully i7nmersed i7t water^ 
it loses weight equal to the weight of a7i equal vohune of water. 
For the proof of the accuracy of this principle, as well as for 
further details about the methods of finding specific gravities, 
a book on Physics must be consulted. 

The specific gravity of brass borings, for example, may be 
found by first weighing them in air and then putting them in the 
specific gravity bottle, filling up with water and again weighing. 

The second weighing will evidently be ; weight of brass + 
weight of bottle + weight of water required to fill empty bottle 
(=50 grams) 7iiinus weight of water displaced by the brass. 
From this we get the weight of an equal bulk of water. In 
all cases the specific gravity is then calculated by dividing the 
weight of the substance by the weight of a7t equal bulk of 
water ^ or 

Weight of known volume 



Specific gravity : 



Weight of equal vol. of water. 



What we have learnt 

In our thirteenth Lesson we have seen that water expands when it is 
heated, and contracts when it is cooled, but that there is a limit to its 
contraction, viz. at 4°, below which temperature further cooling causes it 
to expand until 0° is reached, when a further sudden expansion takes 
place, amounting to about ten per cent of its volume, the water at 0° freezing 
to ice at 0°. Water-pipes are burst by this sudden expansion on freezing, 
and not by the subsequent thaw as is often supposed. 

If we apply heat to ice at say — 10°, its temperature is not raised above 
0°, at which point the application of more heat merely causes it to change 
from the solid to the liquid state. When all the ice is melted, further heat 
raises its temperature until 100° is reached, when the water boils, but 
further application of heat cannot raise its temperature above loo"^, it merely 
causes the water to change its state from a liquid to a gas. 



LESSON XIII SUMMARY AND EXERCISES 115 

Thus in heating ice from — 10° to 100°, we have two stationary points 
as regards temperature, viz. 0° and 100^^, called the freezing point (or 
melting point) and boiling point. The boiling point of water is greatly 
influenced by pressure, water boils at a lower temperature under dimin- 
ished pressure and at a higher temperature under increased pressure. 
The heat absorbed by a body changing its state from solid to liquid or 
liquid to gas, is called latent heat, and is again given out when the body 
changes its state in the opposite direction, viz. from gas to liquid or liquid 
to solid. We have seen how to measure the amount of the latent heat of 
steam, and the latent heat of fusion of ice, and considered the part played 
by latent heat in freezing machines and freezing mixtures. We have 
defined specific heat, and considered Dulong and Petit's Law, that the 
product of the specific heat into the atomic weight of the solid elements is 
approximately constant. We have seen that water is taken as the standard 
with which all liquid and solid bodies are compared as regards their 
density or specific gravity ; and we have seen the relation between specific 
gravity and weight in the metric system, and also learnt how specific gravity 
is found by experiment. 



Exercises on Lesson XIII 

1. Describe the behaviour of water which is (a) cooled from 100° to 
— 10^ ; (d) heated from — 10^ to 100°, mentioning all the important facts 
observed. 

2. What is the influence of pressure on the melting and boiling points 
of water ? 

3. What is meant by latent heat ? How may it be measured ? 

4. Define specific heat and explain Dulong and Petit's Law. 

5. Describe experiments showing how water may be frozen by its own 
evaporation. 

6. Explain the principle of freezing mixtures. 

7. Define a thermal unit. 

8. Define specific gravity. An empty specific gravity bottle w^eighs 
15.4268 grams ; filled with water at 60"^ F. it weighs 66.0694 grams ; when 
filled with a sample of sulphuric acid it weighs 106.2378 grams. What is 
the density of the acid ? 

9. The above bottle weighed 73.4586 grams when 8.4204 grams of brass 
turnings had been placed in it and the bottle filled up with distilled water. 
What is the density of the brass ? 

10. A piece of glass rod weighed in air 4.2882 grams, it weighed in 
water 2.4787 grams. What is its density ? 

11. What weight of ether of specific gravity 0.7204 will the bottle hold 
which is mentioned in No. 8 ? 

12. What is the weight of an iron casting — its volume is 476 c.c. and its 
density 7.436 ? 

13. An iron bottle has a capacity of 784 c.c. What weight of mercury 
will it hold, the density of mercury being 13.59 ? 



LESSON XIV 

WATER AS A SOLVENT — WATER OF CRYSTALLISATION, 
EFFLORESCENCE, DELIQUESCENCE — SOLUBILITY OF 
GASES — NATURAL WATERS — TEMPORARY AND PER- 
MANENT HARDNESS AND THE SOFTENING OF WATER 
— DISTILLATION AND PURIFICATION 

Water as a Solvent. — Water is the most generally use- 
ful of all the known solvents ; nearly all the re-agents of the 
Chemical Laboratory are used in aqueous solution. Not only 
do very many solid substances, especially metallic salts, dis- 
solve in water, but certain liquid compounds, e.g. alcohol and 
acetic acid, as well as nearly all gases, dissolve in water to a 
greater or less extent. 

Solubility of Salts. — The weight of solid compounds 
which can be dissolved in a given weight of water depends 
not only on the nature of the compound, but also on the 
temperature at which solution is effected. The maximum 
amount of each substance which water is capable of dissolving 
at a given temperature is fixed and definite, and when water 
has taken up this amount it is said to be a saturated solution 
of the substance dissolved. 

ExPT. 51. — In order to show the different solubilities of 
salts let us take 30 grams of powdered potassium chlorate and 
place it in a flask with 100 c.c. of water at 30° C. Let us 
now shake the flask until the water becomes saturated ; we see 
that all the salt has not dissolved. Let us now filter the 
solution through filter paper placed in a funnel and evaporate 
the liquid to dryness over a small flame. We see that a white 

tt6 



LESSON XIV WATER AS A SOLVENT 117 

residue is left, which may be collected and heated in a test 
tube to prove that it is potassium chlorate. We see that 
although all this salt has not been dissolved by cold water, yet 
a portion of it has been dissolved. 

ExPT. 52. — Let us now repeat the experiment, but instead 
of shaking up in the cold we will boil the salt with the w^ater ; 
this time it all dissolves, showing the influence of temperature 
on the amount of a salt which can be dissolved. We notice 
that on cooling the solution a portion of the potassium chlorate 
is deposited as crystals. 

ExPT. 53. — Repeat the Experiment 51, but use cold water 
and magnesium sulphate instead of potassium chlorate, but 
in the same quantities. We notice that all this salt dissolves 
even in the cold. 

ExPT. 54. — If we repeat the Experiment 51 with barium 
sulphate we find that the filtered solution gives no residue 
on evaporation, showing that the compound is quite insoluble ; 
and it is so even in boiling water. 

The above facts are wtII indicated in Fig. 60, which repre- 
sents graphically the effect of (i) the nature of the compound, and 
(2) the temperature at which solution is effected. The vertical 
lines indicate temperatures from o^ to 100° C, whilst the 
horizontal lines indicate the number of grams of the salt which 
can be dissolved by 100 grams of water at each particular 
temperature, e.g. we notice on the curves of solubility for 
potassium nitrate and lead nitrate that at 50° both salts are 
equally soluble, 100 grams of water dissolving 85 grams of 
either compound. The two curves also show that the potas- 
sium salt is less soluble at low temperatures and more soluble 
at higher temperatures than is lead nitrate. 

Very many salts dissolve but very slightly in water, 
especially cold water, e.g. lead chloride, PbCU, calcium 
sulphate, CaSO^, strontium sulphate, SrSO^ (very sparingly 
soluble), silver nitrite, AgNO^, silver sulphate, Ag^SO^, 
lead iodide, Pbl2 ; w^hilst others are practically insoluble in 
water, such as barium sulphate, BaSO^, silver chloride, AgCl, 
lead chromate, PbCrO^. 

Precipitation. — Whenever the solutions of two salts are 
mixed together, and by a double decomposition a difficultly 
soluble or insoluble salt can be produced, then that compound 



CURVES OF SOLUBILITY 



separates from the solution in the solid state, and is said to be 
precipitated. Thus ; — 

Pb (NO3), + 2 KI = Pbl2 + 2 KNO3, 
AgNOo + NaCl = AgCl + NaNOg. 

Crystallisation. — When any salt is dissolved in water, 
say boiling water, there is a certain limit to the quantity of 



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salt above which no more salt can be dissolved. The solution 
is, as has been stated, then said to be saturated. If now the 
solution be cooled, the water is usually no longer able to hold 
all the salt in solution, and a portion of it separates out in the 
solid form, and generally in the form of crystals of definite 
shape. Thus from Fig. 60 we may learn that at 100° a 
saturated solution of potassium chlorate contains 60 parts of 



LESSON XIV WATER OF CRYSTALLISATION 119 

the salt per 100 of water, whilst at 15° (the ordinary temper- 
ature of the air) only 5 parts of the salt can be dissolved. 
Evidently, then, a saturated solution at 100° must deposit 55 
parts of the salt as crystals on being cooled down to 15°. 
Similarly a saturated solution of nitre (potassium nitrate), at yoP 
if cooled to 15° must deposit just twice as much, or no parts 
of the salt as crystals. 

Water of Crystallisation. — Very often when metallic 
salts crystallise from aqueous solution they do so as definite 
chemical compounds with water. This water, which is in the 
solid state and in combination wdth the salt (doubtless combined 
with it in the solution also, before crystallisation), is called 
Water of Crystallisation, e.g. 

Sodium sulphate or Glauber's salts . . . Na2S04 . 10 H2O 

Sodium carbonate or washing soda . . . Na2C03 . 10 H2O 

Sodium borate or borax Na2B407 . 10 H2O 

Calcium sulphate or gypsum .... CaS04.2H20 

Copper sulphate or blue vitriol .... CUSO4.5H2O 

Ferrous sulphate or green vitriol .... FeS04.7H20 

Zinc sulphate or white vitriol .... ZnS04 . 7 H2O 

Magnesium sulphate or Epsom salts . . . MgS04 • 7 H2O 
Aluminium potassium sulphate or 

common alum Al2(S04)3K2S04.24H20 

Efflorescence. — Many salts containing water of crystal- 
hsation are not able to retain it all in a dry atmosphere, and 
on exposure to the air they become opaque, lose a part of 
their water and fall to powder ; they are then said to effloresce. 

ExPT. 55. — Expose a few crystals of washing soda to a 
dry warm atmosphere and notice their gradual efflorescence. 

Deliquescence. — Other salts, on the other hand, attract 
water with such great avidity (when they are said to be hygro- 
scopic) that they liquefy if exposed to the air owing to the 
absorption of moisture, e.g. calcium chloride, potassium acetate, 
and ferric chloride. The first named is used in the anhydrous * 
state as a drying agent. 

ExPT. 56. — Expose some dry calcium chloride to the air in 
a dish and notice how soon it becomes moist and gradually 
deliquesces to a liquid. 

* Anhydrous means without water. 



I20 SOLUBILITY OF GASES part ii 

Solubility of G-ases. — All gases dissolve to a greater or 
less extent in water. 

Such gases as hydrogen, oxygen, nitrogen, air, or nitric 
oxide, dissolve but slightly in water, thus a litre of water will 
dissolve only 20.92 c.c. of oxygen and nitrogen from the air at 
15° C. (see p. 135). Other gases, such as nitrous oxide, chlorine, 
sulphuretted hydrogen, and carbon dioxide, are more soluble, 
whereas other gases, such as hydrochloric acid, HCl, ammonia, 
NH3, sulphur dioxide, SO2, are exceedingly soluble in water. 

Influence of Temperature. — The effect of temperature 
on the solubility of gases in water is the opposite to that on 
the solubility of solids, for the higher the temperature the less 
gas is dissolved, and boiled water, especially if boiled in a 
vacuum, contains no dissolved gases ; e.g. a litre of w^ater 
dissolves 

23.34 c.c. of nitrogen and oxygen from air at 10^ C. 
19.20 " " " 20^ C. 

17.69 " " " 25° C. 

Influence of Pressure. (Law of Henry). — The volume 
of gas (measured at a constant pressure) which can be dis- 
solved by w^ater is proportional to the pressure at which the 
solution is effected. Thus a litre of water will absorb 4 
times as much carbon dioxide under a pressure of 4 atmo- 
spheres as it would under a pressure of i atmosphere. This 
law was expressed by the discoverer (Henry) that ^^ under 
equal circumstances of temperature, water takes up in all 
cases the same volu?ne of coiidensed gas as of gas inider 
ordiiiary pressure.'^'' This is evidently the same as the first 
statement, for from Boyle's Law we know that the volume of a 
gas is inversely proportional to the pressure. Therefore 4 
volumes at i atmosphere become i volume at 4 atmospheres, 
and if i volume of gas is dissolved under i atmosphere, under 
4 atmospheres i volume of gas wall still be dissolved ; but as 
this at consta7it pressure (i atmosphere) would be 4 volumes, 
we may say that the volume dissolved (measured at constant 
pressure) is proportional to the pressure. We can now under- 
stand why it is that on opening a bottle of soda-water the 
liquid begins to effervesce. It is because carbon dioxide has 



LESSON XIV DISSOLVED AIR 



been dissolved under pressure^ and has taken up much more 
than it can hold under the ordinary pressure of the air, hence 
a portion of it escapes immediately the pressure is relieved by 
opening the bottle. 

ExPT. 57. — Take a soda-water bottle which is closed by 
an ordinary cork, and pierce the cork with a hollow borer 
which is connected with a piece of tubing passing under the 
bee-hive shelf of the pneumatic trough. When the cork is 
quite pierced a sudden rush of carbonic acid will be noticed, 
and this can be collected in the gas jar placed over water to 
receive it. 

Use of dissolved Air in Natural Waters. — The dis- 
solved air in water is of great importance in nature. By its 
means fishes and all aquatic animals are supplied with the oxygen 
necessary for their life. In water which has been boiled, so as 
to expel the dissolved air, and then allowed to cool, fishes soon 
die for want of air. In water analysis also, the determination 
of the amount of dissolved oxygen is of importance, because 
if a particular water is very impure, minute living organisms 
called bacteria thrive in it, and these use up the dissolved 
oxygen ; this loss of free oxygen can be detected by analysis, 
and so an inference drawn as to presence or absence of the 
organic impurities which favour the growth of the bacteria and 
the consequent loss of dissolved oxygen. Boiled water which 
has been allowed to cool in absence of air is insipid and taste- 
less, whereas a glass of spring water is refreshing, a difference 
due almost entirely to the presence or absence of dissolved air. 

Natural Waters. — We may divide natural waters into 
(i) Rain Water, (2) Spring Water, (3) River Water, and (4) 
Sea Water. 

Rain Water. — Rain water is the purest form of natural 
water, and may indeed be called natural distilled water, for it 
is produced by the slow evaporation (or distillation) of the 
water on the surface of the earth, the vapour of which rising 
into colder regions of the atmosphere becomes condensed in 
the form of clouds, which on further cooling fall as rain. 
Although rain water is free from the soluble impurities derived 
from the surface of the ground, yet it is not perfectly pure, 
because, in falling through our atmosphere, it has washed out 
certain impurities which are found there in larger or smaller 



NATURAL WATERS PART II 



quantities, e.g. near large towns the atmosphere contains small 
quantities of tarry matter, soot, as well as sulphur dioxide and sul- 
phuric acid, both derived from the burning of coal which contains 
small quantities of iron pyrites (FeS^), hence such water has an 
acid reaction. Rain w^ater also dissolves a certain amount of 
free nitrogen and ox3'gen as well as carbon dioxide from the 
air (see p. 135). It contains, in addition, minute quantities of 
common salt, ammoniacal salts, nitrites and nitrates, and 
organic matter as well as mineral particles which float in the 
atmosphere as dust. 

Spring Water. — When rain w^ater percolates through the 
earth it dissolves a certain amount of the materials with which it 
comes in contact ; the quantity and kind of impurity thus taken 
up depends upon the nature of the soil and strata through 
which the water passes. Hence when the ^vater emerges at 
the surface of the earth as a spring at a lower level than that 
of the gathering ground, it always contains dissolved mineral 
matter as wxll as a larger proportion of carbon dioxide derived 
from the soil. 

River Water. — River water differs from spring w^ater in 
its not having, as a rule, filtered through various depths of 
porous strata. It is consequently generally turbid, and con- 
tains more or less insoluble matter in suspension, whereas 
spring W'ater is usually clear and sparkling. River w^ater may 
be greatly purified by filtration through beds of sand or other 
material, but this process only removes the suspended impu- 
rities. The clear filtered water still contains soluble impurities 
just as spring water does. On a large scale water is sometimes 
softened by the addition of lime before it is filtered. The 
composition of the mineral impurities of river w^ater varies 
considerably wdth the nature of the ground over which the 
water runs; thus Thames water contains about 11 grains per 
gallon of carbonate of lime, the Trent 21 grains of sulphate 
of lime. They are both hard waters ; * the first is temporarily 
hard, whilst the second is permanently hard. The waters of 
the Dee and the Don in Aberdeenshire, draining a granitic 
district, are, on the other hand, soft waters. In addition to 
their natural ingredients river water is very often polluted with 
the sewage of towns and the refuse from manufactories and 
* For an explanation of hard and soft waters see p. 123. 



I 



LESSON XIV HARD AND SOFT WATERS 123 

mines, and it is now an important matter to prevent this pollu- 
tion, and to devise suitable means of purifying such liquids 
before they flow into the rivers. 

Sea Water. — Sea v^ater contains the largest amount of 
dissolved solids, the mean quantity being about 3.6 per cent, of 
which common salt (NaCl) makes up 2.6 per cent, the rest 
being mainly composed of the sulphates, chlorides, and carbon- 
ates of magnesia, lime, and potash. 

Hard and Soft Waters. — When the mineral matters in 
solution in natural waters have no action on soap and do not 
retard the formation of a lather when one washes the hands, 
such water is said to be soft water. Rain water is of course 
soft water, because it has had no opportunity of dissolving 
saline matter, so also is water which has only flowed over 
granite rocks or sandstone, which are insoluble in water. But 
rain water which has come in contact with limestone rocks or 
chalk (made up mainly of calcium carbonate, CaCOg), or which 
contains gypsum (calcium sulphate, CaSO^) in solution, decom- 
poses a portion of the soap and so retards the formation of a 
lather in washing. Hard water renders a portion of the soap 
useless, because it produces an insoluble compound with the 
soap. Hardness in water may be of two kinds, either (i) tem- 
porary hardness, or (2) permanent hardness. 

Temporary Hardness. — The presence of dissolved carbon 
dioxide in water has an important effect as regards its power of 
dissolving certain mineral substances. 

ExPT. 58. — If we take a dilute solution of lime-water and 
pass a current of carbon dioxide through it, we notice first that 
a white precipitate of calcium carbonate is formed, for this 
substance is insoluble in pure water. 

Ca(0H)2 + CO2 = CaCO, + H^O. 

Lime-water. Calcium carbonate. 

If, however, we continue to pass the gas through the turbid 
liquid we notice that it gradually becomes clear, and soon all 
the precipitate disappears ; this is because the water now con- 
tains free carbonic acid ; thus : — 

H,0 + CO^ = H2CO,. 

Carbon dioxide. Carbonic acid. 



124 • PERMANENT HARDNESS PART ii 

This has the power of dissolving calcium carbonate, the bi- 
carbonate * of calcium being formed ; thus : — 

CaCOg + H2CO3 = Ca(HC03)2. 

Normal calcium carbonate. Calcium bicarbonate. 

Bicarbonate of lime is the principal substance which causes 
temporary hardness in water, and it may be taken as a type of 
the rest. This hardness is called, temporary because it may be 
got rid of by boiling, or by precipitation with lime. By this 
means the bicarbonate is decomposed, the normal carbonate 
being precipitated, whilst carbonic acid is freed and expelled as 
CO., ; thus : — 

Ca(HC0o)2 = CaCOg + H^O + CO^. 

ExPT. 59. — Boil a portion of the clear liquid obtained by 
passing CO^ into lime-water. A white precipitate is again 
formed, whilst bubbles of CO2 are given off. 

If we shake up equal volumes of the boiled and unboiled 
water with a measured volume of soap solution added gradually, 
we shall find that a much larger quantity of soap is required to 
produce a lather in the case of the unboiled or hard water than 
is required for the boiled or softened water. 

Clark's Process for softening Water. — Temporary 
hardness due to calcium bicarbonate may be got rid of and 
the water softened by the addition of a proper proportion of 
lime, sufficient in quantity to combine with the free carbonic 
acid and also to decompose the bicarbonate of lime present ; 
thus : — 

Ca(HC03)2 + Ca(0H)2 = 2 CaCO.. + 2 H2O. 

The result is that the normal carbonate of lime is produced in 
a liquid which contains no free carbonic acid to dissolve it, 
hence it falls as an insoluble precipitate. f 

Permanent Hardness. — Hardness w^hich cannot thus be 
got rid of by boiling is said to be permanent hardness. It is 
mainly due to gypsum, CaSO^, or to the sulphates or chlorides 

* The student should refer to page 205 for an explanation of normal and 
acid salts. 

t Temporary hardness may be due to the presence of the bicarbonates 
of lime, magnesia, iron, or manganese. 



LESSON XIV PURIFICATION OF WATER 125 

of lime and magnesia which are sokible in water. Boiling has 
no effect on the solutions of these salts, but they may be decom- 
posed by the addition of common washing soda, sodium carbon- 
ate, NagCOo, 10 H^O ; thus : — 

CaSO^ + Na^COg = CaCOo + Na^SO^. 

Hence the use of washing soda for softening water. The sodium 
sulphate produced by the double decomposition has no effect on 
the soap. Water softened by the addition of sodium carbon- 
ate is not suited for drinking purposes, inasmuch as it contains 
sodium sulphate in solution (see equation). 

ExPT. 60. — Make a dilute solution of calcium sulphate and 
divide it into two portions ; boil one of them and observe that 
there is no precipitate as in the former case. Now test both 
waters with soap as before ; both are alike, showing that boiling 
has had no effect. Soften a third portion of the solution by the 
addition of a little washing soda, and test again with soap solu- 
tion. It will be seen that the hardness has disappeared. 

Purification of Natural Waters. — Natural waters, espe- 
cially river waters (e.g. the Thames), although not pure enough 
for domestic or drinking purposes, may be rendered fit for such 
use by a proper method of purification. We may divide such 
processes into (i) softening, (2) precipitation or subsidence, 
(3) filtration. 

The method of softening hard waters has already been men- 
tioned, on the large scale ; softening by lime is usually employed 
(Clark's process), and the precipitated carbonate of lime allowed 
to subside. This subsidence of solid particles or precipitation 
carries down a large proportion of the suspended matter, the 
remainder being got rid of by filtration through large filter-beds 
of sand and gravel. Subsidence and filtration only free the 
water from suspended matter, whilst the chemical process of 
softening removes the soluble bicarbonates and sulphates of 
lime, magnesia, etc. The other soluble impurities cannot be got 
rid of by these methods. 

Distillation of Water. — In order to obtain water free from 
these soluble impurities it must be distilled. 

ExPT. 61. — Place some water coloured blue with copper sul- 
phate in a retort (Fig. 61), and heat it with a Bunsen burner. 



126 



DISTILLATION OF WATER 



PART II 



When it begins to boil and the steam passes through the glass 
tube, which is cooled on the outside by a jacket through which 
a current of cold water flows, it will be seen that the condensed 
steam is quite colourless water, and that it has been separated 
from the soluble impurity (copper sulphate) which is left behind 
in the still. 

Volatile Impurities in Water. — In order to obtain pure 
water, however, simple distillation is not sufficient, as some of 
the impurities derived from organic matter are volatile, and pass 
over with the steam into the condenser, and so obtain access to 
the distilled water. 

ExPT. 62. — Repeat Experiment 61 with the addition of a sin- 
gle drop of ammonia instead of copper sulphate. The ammonia 




Fig. 61. 



may be detected in the clear distillate by adding i c.c. of 
Nessler's solution,* when a deep-brown colouration is produced, 
showing that even the distilled water contains a volatile im- 
purity. 

In order to obtain water free from these volatile impurities, 
the water must first be boiled with a strong solution of potas- 
sium permanganate, KMnO^, and caustic potash, KOH. These 
oxidise and decompose the organic matter, and the volatile 
products pass over into the condenser with the first portion 
of steam, the condensed water from which must be rejected as 

* Nessler's solution is made by dissolving 62.5 grams of potassium 
iodide in about 250 c.c. of water, and gradually adding a cold saturated 
solution of mercuric chloride until a slight permanent precipitate is pro- 
duced. Then 150 grams of caustic potash is added, and the solution made 
up to a litre in volume. 



LESSON XIV SUMMARY AND EXERCISES 127 

impure. When about one-third of the water has been distilled, 
the middle third may be collected, and is free from both soluble 
and volatile impurities if proper precautions are taken in the 
distillation. 

Even distilled water becomes impure if freely exposed to 
the air, for it dissolves the atmospheric gases and traces of 
the other impurities in the atmosphere. Pure distilled water 
is now made on a large scale, and sold to the public under 
various names after aeration with carbon dioxide. 



What we have learnt 

In our fourteenth Lesson we have studied water as a solvent, and noticed 
the influence of temperature on the amount of a soUd which can be dis- 
solved in water. It was noticed that solids are usually more soluble at 
higher than at lower temperatures, whilst gases are less soluble. 

When the solutions of two salts are mixed together, precipitation takes 
place, when they can produce an insoluble or difficultly soluble compound 
by their mutual decomposition. Many metallic salts combine with water, 
and the compounds thus produced are capable of existing in the solid state 
as crystals. We have seen that the volume of a gas (measured at a constant 
pressure) which can be dissolved by water is proportional to the pressure 
at which the solution is effected. The air dissolved in water is of the 
utmost importance for the preservation of all forms of aquatic animal life. 
The most important characters of rain water, spring, river, and sea-water 
have been mentioned, whilst the causes of temporary and permanent hard- 
ness in water have been explained. The former being due principally to 
the presence of the bicarbonates of lime and magnesia, and the latter to 
their sulphates or chlorides. 

It was seen that temporary hardness could be removed by boiling, or 
by precipitation with Hme (Clark's process), whilst permanent hardness 
could not be removed by these means, but could be removed by the addi- 
tion of sodium carbonate. 

The various methods of purifying water from its suspended and dissolved 
impurities such as subsidence, softening, precipitation, filtration, and dis- 
tillation have been explained. 



. Exercises on Lesson XIV 

1. What is the effect of temperature on the solubility of (i) a solid, 
(2) a gas ? What is the effect of pressure on the solubility of a gas ? 

2. Explain the precipitation which often takes place when the solution 
of two salts are mixed together. 



128 EXERCISES 



3. Give the formulae and common names of several salts which contain 
water of crystallisation. 

4. What are the chemical characters of the various natural waters ? 

5. Of what importance is the dissolved air in natural waters ? 

6. Explain the hardness of water, both temporary and permanent, state 
how such waters may be softened, and give equations for the reactions. 

7. Explain fully how you would prepare pure water from a muddy river 
water. 

8. Is distilled water always pure ? If not, why not ? 

9. What is meant by efflorescence and deliquescence ? 

10. What weight of anhydrous sedium carbonate is contained in 1000 
grams of washing-soda ? 



LESSON XV 

NITROGEN AND AIR 

PRESSURE, TEMPERATURE, HUMIDITY, AND EXTENT OF 
THE ATMOSPHERE — THE BAROMETER — CHEMICAL 
COMPOSITION AND ANALYSIS OF AIR — ACTION OF 
ANIMALS AND PLANTS ON THE AIR — VENTILATION 

Nitrogen 

Symbol N. Atomic Weight 13.94. Density 13.94 

Occurrence. — As we have seen (p. 9) nitrogen occurs in 
the free state in the air, of which it forms four-fifths by volume. 
Nitrogen exists in the state of combination in nitre or salt- 
petre, potassium nitrate, KNO3 (whence the element derives 
its name). It forms an essential constituent of the bodies of 
all animals and plants. 

Preparation. — (i) Nitrogen can be prepared (see Expt. 9) 
by burning phosphorus in a closed volume of air, the oxygen 
being entirely removed ; after the white fumes of phosphorus 
pentoxide (P2O5) have disappeared, the colourless gas remain- 
ing is nearly pure nitrogen (see Fig. 62) . 

(2) To obtain it perfectly pure, air which has been dried and 
freed from carbonic acid, is passed over metallic copper turnings 
in a long glass tube heated to redness, when oxide of copper 
remains in the tube and pure nitrogen passes over and may be 
collected in a pneumatic trough. 

(3) Nitrogen gas can also be obtained by heating certain 
of its compounds. Thus if a concentrated solution of am- 

129 



I30 



NITROGEN 



monium nitrite (NH^NOg) be heated, nitrogen is evolved 
whilst water is eliminated. 



(NHJNO, 



N2 + 2H2O. 



(4) Ammonia, a compound of nitrogen and hydrogen (NH3), 
is decomposed by chlorine, with formation of hydrochloric acid 
and liberation of nitrogen : — 

2 NH3 + 3 CI2 = 6 HCl + N2. 

If a current of chlorine gas be passed through a satu- 
rated solution of ammonia the re- 
action as above stated occurs, and 
nitrogen can be collected as usual. 
The hydrochloric acid formed, unites 
wdth the excess of ammonia form- 
ing salammoniac (NH^Cl). 

8 NHg + 3 Cl^ = No + 6 NH4CI. 

II I ^ '" III Care must be taken, in mak- 

ing this experiment, that ammonia 
always remains in excess, other- 
llj wise an explosive compound is pro- 

duced. 

Properties. — Although nitro- 
gen does not support combustion 
and is an inert gas, yet it can be 
made to combine with both oxygen 
and hydrogen ; with both elements 
it unites to form a strong acid, viz. nitric acid, HNOo, and 
with the latter it combines to form a pow^erful base, ammonia, 
NH3. At a very low^ temperature and under a high pressure, 
nitrogen condenses to a colourless liquid which boils about 
— 193°. Nitrogen acts as a diluent to the atmospheric oxygen, 
and although it does not support life, it is not poisonous, 
animals brought into nitrogen dying of suffocation from want 
of oxygen ; indeed, it is clear that nitrogen cannot exert an 
injurious effect on the animal as it is inhaled in large quantities 
at every inspiration. 




Fig. 62. 



LESSON XV THE ATMOSPHERE 131 



NITROGEN AND OXYGEN 
The Atmosphere 

Experiments described in the foregoing lessons have 
shown us that the air or atmosphere (ar/xo? vapour, crc^atpa 
a sphere) which surrounds the earth consists of oxygen and 
nitrogen gases present in the proportion of about i to 4 by 
volume. It now becomes necessary to study the properties 
as well as the composition of the air more closely, and in the 
first place the question of the weight of the air claims 
attention. 

Weight of the Air. — This can readily be, shown by 
removing the air, by means of an air-pump, from a glass 
globe, the neck of which is furnished with a stopcock. 
Hanging the vacuous globe on one arm of a balance, and 
placing weights to equipoise the system in the pan attached 
to the other arm. On then opening the stopcock, the air 
will be heard to rush in and the equilibrium will be destroyed, 
the arm of the balance to which the globe is attached descend- 
ing. The weight of the air has been very exactly determined 
by a process similar to that above described, but carried out 
with every precaution, so that the errors of experiment are 
reduced to as small an amount as possible. These determina- 
tions show that one litre of dry air at 0° C. and under a 
pressure of 760 mm. at the sea's level at London (latitude 
51° 28') weighs 1.293 grams. 

Pressure of the Atmosphere. — The atmosphere thus 
possessing weight, must exert a pressure on every object on 
the earth's surface, and therefore upon the human body. But 
inasmuch as the pressure of the air is exerted, under ordinary 
conditions, in all directions equally, we do not feel it. If^ 
however, we place the hand so as to close the top of a glass 
cylinder, open at both ends, setting the other end firmly on 
the plate of an air-pump which is then w^orked so as to 
withdraw the air from within the cylinder, the fact of the 
weight or pressure of the air becomes very evident and we 
shall feel some difficulty in removing the hand. If a piece of 
bladder be tied over the mouth of the cylinder instead of the 



132 THE BAROMETER part ii 

hand being placed on it, it will be found that as the air 
is abstracted from beneath, the bladder bulges inwards and 
at last bursts with a loud noise owing to the pressure of air 
upon it. As the air is an elastic fluid, it obeys Boyle's law 
of pressures (see p. 53), that is, its density is directly pro- 
portional to the pressure to which it is subjected. Hence, 
as we ascend from the sea^s level, the density of the air 
becomes less, a portion of the superincumbent pressure 
being removed. At the sea^s level the average pressure of 
the atmosphere is equal to that of a column of mercury of 760 
mm. or 29.922 inches, and as mercury is 13.5 times as heavy 
as water, the air can sustain a column of water of 32 feet. 
This pressure amounts to 103.3 kilos on every square centi- 
metre or 14.73 lbs. per square inch. The human body, has, 
therefore, to sustain a pressure of several tons, and if this 
pressure be suddenly reduced, as when rapid ascents are 
made in balloons, the effect of the diminution is observed in 
haemorrhage from the nose, mouth, and eyes, consequent upon 
the bursting of small blood-vessels in those parts. The same 
elevation may be attained by gradual means, as by ascending 
a mountain, without evil effects resulting, as the blood-vessels 
then have time to accommodate themselves to the new con- 
ditions. 

The Barometer. — In the seventeenth century a Florentine 
pump-maker found that an ordinary lift-pump would not raise 
water more than 32 feet. In 1643 Torricelli gave the true 
explanation by making a simple experiment. Taking a glass 
tube 3 feet in length, closed at one end, and filling it with 
mercury he closed the open end with his thumb, and then 
inverted the tube in a basin filled with mercury. The mer- 
cury then sank in the tube till the column was about 760 mm. 
in height, at which point it remained stationary. Above 
this level was an empty space which is still called the Tor- 
ricellian vacuum. He next poured water on to the top of 
the mercury contained in the basin, and raised the tube so 
that the open end came into the water. The mercury then 
flowed out and the water rushed up, completely filling the tube. 
Thus the baro?neter was discovered, but this name was given 
to it by Robert Boyle. In order to test the accuracy of the 
explanation of the suspension of the mercury in the tube 



LESSON XV 



ATMOSPHERIC PRESSURE 



133 



being due to the atmospheric pressure, Pascal determined to 
ascertain whether the mercury in the barometer sank when the 
instrument was taken up a mountain. Unable to do this himself 
he commissioned his brother-in-law Perier to make the trial, and 
this he did on the Puy-du-D6me. The mercury continued to 
sink as he ascended the mountain until he reached the summit, 
when it remained at a constant level. 

The sinking of the barometric column as the superincumbent 
pressure is removed is shown by a simple experiment. Here I 
have a barometer-tube (Fig. 63) filled with mercury; I invert 
it in a basin of mercury and we see that the mercury sinks to 
a certain level, the space above being a 
vacuum. Now a tubulated receiver is 
brought over the tube, a perforated india- 
rubber stopper surrounding the tube and 
fitting tightly into the tubulus of the re- 
ceiver, the whole being placed on the plate 
of an air-pump. As the air in the receiver 
is pumped out, we notice that the level of 
the mercury in the tube gradually sinks, 
until at last it is nearly on the same level as 
the mercury in the basin. On slowly open- 
ing the stopcock of the pump, air rushes in 
to fill the empty receiver, and the mercury 
in the tube is seen gradually to rise until it 
reaches the point at which it stood before 
the experiment. 

The variation in the height of the baro- 
meter with elevation above the sea's level can be employed to de- 
termine the amount of that elevation. When the mercury sinks 
one-tenth of an inch, the elevation reached is about 90 feet, but 
for this purpose the cumbersome mercurial barometer is generally 
discarded in favour of a little pocket aneroid barometer which is 
about as large as an ordinary watch, and in which no mercury is 
used. The variations in pressure cause an index to move over a 
graduated dial. In making mercurial barometers care must be 
taken to exclude all air or other gas from the tube, as the pres- 
sure of even a trace greatly impairs the instrument. It is there- 
fore necessary to boil the mercury in the tube so as to expel all 
air adhering to the glass or to the mercury. 




Fig. 63. 



134 EXTENT OF THE ATMOSPHERE part ii 

Variations in Height of the Barometer. — As the atmo- 
sphere is in constant motion, its pressure at the same spot un- 
dergoes changes, and these are indicated by corresponding 
changes in the height of the barometer, and as a change of 
atmospheric pressure usually accompanies change of weather, 
the variations of the barometer indicate approaching change of 
meteorological conditions. Hence the use of the barometer as a 
weather glass. 

Extent of the Atmosphere. — As the density of the air 
diminishes as the distance from the sea^s level increases, it is 
difficult to say exactly to what height it extends. But there is 
no doubt that a limit exists beyond which there is no air. 
The height at which the air still possesses sensible density is 
calculated to be from 40 to 45 miles. Suppose the earth to be 
represented by a globe one foot in diameter, the atmosphere 
will be represented by a film of air jL of an inch thick. If 
the air were of the same density throughout, it would only 
reach to a height of rather more than 5 miles above the sea's 
level. 

Temperature of the Air. — Like its density, the tempera- 
ture of the air varies greatly at different places, and at different 
heights over the same place. It diminishes everywhere as we 
ascend into higher regions, and at last we come to a point where 
the temperature never rises above the freezing point. This is 
called the line of perpetual snow. Under the equator this line is 
not reached until we rise above 15,000 feet; in our own latitude 
it is reached at a height of about 4000 feet, whilst in the latitude 
of 75"^ it comes down to the sea's level. The atmosphere acts 
not only as a powerful modifier of the sun's heat, but as an 
equally powerful retainer and distributer of that heat. Were 
it not for the presence of the air, the earth's surface would be 
scorched by day and frozen by night. 

Air a Mixture, not a Compound. — In the first place let 
us discuss the question. Is the air a compound of oxygen and 
nitrogen, or a mere mechanical mixture ? For some time chem- 
ists were doubtful what answer to give. Now, there is no doubt 
that air is a mechanical mixture, and not a chemical compound, 
and for the following reasons : — 

(i) When gases combine heat is given out, and often an 
alteration of bulk takes place. If we mix oxygen and nitrogen 



LESSON XV AIR A MIXTURE OF GASES 135 

in the proportion in which they occur in air, no evolution of heat 
or change of bulk can be noticed^ and the mixture acts in every 
respect like air. 

(2) Gases combine in the ratio of their atomic weights, or 
in multiples of these ; the ratio between oxygen and nitrogen 
in the air is not that of the combining weights of oxygen and 
nitrogen, or in multiples thereof. 

(3) Although the composition of the air is nearly constant 
as regards the proportion of oxygen to nitrogen, yet it is not 
exactly so, small but perceptible variations occur. No such 
variations could be observed if these gases wxre combined in 
air to form a chemical combination. 

(4) And most convincing proof. If we shake up air with 
water, the water absorbs some of the air ; if we take the water 
thus saturated with air, and boil it, the dissolved air will escape. 
The air which thus escapes can be collected and analysed. 
This has often been done, and it is found that the quantity of 
oxygen contained in the air which has been dissolved by the 
w^ater, is in larger proportion to the nitrogen in the same dis- 
solved air than it was in the original atmospheric air. Thus, 
whilst the percentage of oxygen in the air taken is nearly 21, 
that in the air obtained by boiling the water is nearly 34. 





Air undissolved 
IN Water. 


Air dissolved 
IN Water. 


Oxygen 

Nitrogen . 


20.96 
79-04 


33-64 
66.36 




100.00 


100.00 



This change in the proportion of nitrogen to oxygen from 
4:1 to nearly 2:1, could not occur if the air were a com- 
pound, as the compound would be dissolved as a %vhole ; it 
would not be decomposed by simply shaking it up with water, 
and then the relative proportion between the tw^o gases would 
remain, after boiling out, the same as it was before, viz., 4 to i. 

The above numbers exactly agree with the solubilities of 
oxygen and nitrogen separately. Oxygen is found to be more 
soluble than nitrogen. 

(5) Another convincing proof is that when air is liquefied 
by great cold and pressure, it is found that gaseous nitrogen 
is first given off from the liquid, and after a while, the gas 



CO 



136 



ANALYSIS OF AIR 



PART II 



which is slowly evolved, contains sufficient oxygen to ignite a 
glowdng chip. Thus the liquefied air behaves as a mechanical 
mixture of the two elements, one of which (nitrogen) is more 
volatile than the other. Hence we have five reasons for regard- 
ing air as a mechanical mixture of its two chief constitutents. 

How Air is Analysed. — If we wish to determine with 
accuracy the composition of the atmosphere, we must use more 
exact methods than those described in the foregoing lessons. 
First, as regards oxygen and nitrogen. Chemists use for this 
purpose two methods. One to ascertain exactly the weights 
of oxygen and nitrogen present ; the second to measure with 
accuracy the volumes of these two gases present in air; for 




Fig. 64. 



both methods the air must be previously freed from all other 
ingredients. 

Analysis of Air by "Weight. — This method depends on 
the fact that oxygen gas is absorbed by heated metallic copper 
with formation of oxide of copper, CuO, whilst nitrogen gas 
is not thus absorbed. Fig. 64 shows the form of apparatus 
used for this purpose. The glass globe V, furnished wdth a 
stopcock u^ is rendered vacuous by the air-pump, and, after 
being carefully weighed, is attached to a tube of difficultly- 
fusible glass {aU) which is furnished with tw^o stopcocks (rr), 
this is also carefully weighed after having been filled with 
bright copper turnings. This tube is connected at the one end 



LESSON XV GRAVIMETRIC COMPOSITION 137 

with the empty globe, and at the other with a series of tubes 
and bulbs (C, B, and A). These serve to free the air from 
carbonic acid, by means of the bulbs which contain a strong 
solution of caustic soda ; and from aqueous vapour, by means 
of the U tubes (B and C) which contain pumice stone soaked 
in strong sulphuric acid, a substance which absorbs aqueous 
vapour greedily. The weighed tube containing the copper is 
then heated to dull redness by the Bunsen burners fixed in the 
gas furnace. Next, the stopcocks rr and // are slightly opened 
when the air of the room in which the experiment is made, or 
that brought from a distance in a glass globe and attached to 
the bulbs (A), passes over the hot copper, which thereby is 
oxidised, and the air thus completely deprived of its oxygen, 
whilst the nitrogen passes unabsorbed into the globe V. As 
soon as the experiment is finished, and the tube allow^ed to 
cooL both the globe and the cool tube are weighed. The 
increase in weight of the globe gives the weight of the nitrogen, 
wliilst that of the tube gives that of the oxygen. In an experi- 
ment of this kind carried out with very great care, the weight 
of oxygen obtained w^as found to be 3.680 grams ; that of the 
nitrogen being 12.373 grams. Hence the percentage com- 
position of air by weight is: O 23.005 per cent; N 76.995 
per cent ; and many repetitions of this experiment proved that 
the 100 parts by weight of air contains almost exactly 23 parts 
of oxygen to ']^ parts of nitrogen. 

As we know the relative densities of oxygen and nitrogen 
it is easy to calculate from the above numbers the com- 
position of air by vohiine^ and we find that O = 20.77 P^^ cent 
and N = 79.23 per cent. Hence, comparing the two results, 
we have — 

Composition of the Air 

By Volume. By Weight. 
Oxygen .... 20.77 23.005 

Nitrogen .... 79.23 76.995 

100.00 100.00 



Composition of Air by Volume. — But it is necessary 
that WT should have the means of determining the composition 
by volume experimentally. For this purpose, we employ an 



138 EUDIOMETRIC ANALYSIS part ii 

instrument called a Eudiometer (from evSia, clear weather ; and 
ixerpov, a measure (a measure of the clearness or purity of the 
air, or of the quantity of oxygen which it contains). The 
arrangement used is shown in Fig. 46, p. 93. 

The eudiometer A, a strong tube closed at one end, and care- 
fully graduated in millimetres, is first completely filled with 
mercury, and such a volume of the air to be examined is 
allowed to enter as will fill about J of the eudiometer. This 
volume is then accurately read off by a telescope on the divided 
millimetre scale of the eudiometer, the height of the barometer 
and that of the thermometer as well as that of the column 
of mercury in the eudiometer above the level of the mercury 
in the trough being also ascertained for the reasons given 
on p. 93. Next, a volume of pure hydrogen, more than 
enough to combine with all the oxygen contained in the air 
taken, is added, and the volume of air plus hydrogen, carefully 
ascertained with all the precautions already mentioned. The 
open end of the eudiometer is then firmly held on a thick 
sheet of india-rubber under the mercury, and an electric spark 
passed through the gases by means of a battery and induction 
coil, which passes the spark through the gas by two platinum 
wires fused into the sides of the tube at its upper end. Com- 
bination of the hydrogen with the oxygen of the air occurs in 
the eudiometer, water being formed. The eudiometer is 
then raised from the india-rubber block, when the mercury 
is seen to rise in the tube, showing that a diminution of 
volume has taken place, this is accurately measured as before. 
To what is this due? It is caused, as we have learnt 
(p. 92), by the union of exactly two volumes of hydrogen 
with one volume of oxygen, to form liquid water, which 
occupies so small a space, compared with that of the con- 
stituent gases, as to be inappreciable. Therefore, the volume 
of oxygen contained in the air, under examination, must be 
one-third of the diminution in volume which has been obser^xd 
as a consequence of the combination. If I took exactly 100 
volumes of air, and added 50 volumes of hydrogen, and I found 
87 volumes after the explosion, or a diminution of 6^ volumes, 

then — = 21 w'ould be the volume of oxygen contained in 100 

volumes of air, the other 79 being nitrogen. 





428.93 






74977 






480,09 






269.68 




tion 


89.89 


= 20.96 




339.04 -- 


= 79.04 N 
100.00 



LESSON XV CARBON DIOXIDE IN AIR 139 



Example 

Air employed (vols, corrected to NTP) 

Air + hydrogen 

After the explosion 

Diminution in volume ( = water H.2O) 

Oxygen corresponding to diminution = ^ diminui 

Air — oxygen = nitrogen 



Results of Air Analyses as regards Oxygen and 
Nitrogen. — A large number of eudiometric analyses of air 
made at different times at the same locality, and at different 
localities, have shown that although the proportion of 21 
vohjmes of oxygen to 79 of nitrogen may be taken as giving the 
average composition of pure air, yet differences are observed 
which lie outside the errors of experiment. The variation 
ranges from 20.9 to 21.0, and in the air of towns, and espe- 
cially in mines, the amount of oxygen may sink as low as 20.3. 
The low percentage of oxygen 20.77 found by the gravimetric 
method must be taken as being not so accurate as that obtained 
by the volumetric method 20.96, as there are slight unavoidable 
experimental errors in the former, which are got rid of in the 
latter method. 

ExpT. 63. Carbonic Acid in the Air. — If lime-water 
or baryta water * is poured into a shallow dish, and exposed to 
the air for half an hour, it will be noticed, on pouring it into 
a test glass, that it has become turbid, and that white particles 
are floating about in the liquid. As this is the test we pre- 
viously used to detect carbonic acid gas (or carbon dioxide) 
the experiment shows that this gas is present in atmospheric 
air, and we have already learnt that fires and the breathing of 
animals produce it. It is only present, however, in minute 

* Lime-water is made from quicklime, CaO, by pouring water over it so 
as to slake it, in which operation heat is evolved, owing to the chemical 
combination between the lime and water. CaO + H2O = Ca(OH)2. Slaked 
lime or calcium hydrate, Ca(OH)2, dissolves slightly in water, and a clear 
solution is obtained on shaking up some of the hydrate in water and allow- 
ing the excess to subside. The clear solution may be poured off after a 
few hours, and kept in a well-stoppered bottle for use. Baryta water is 
similarly prepared. 



140 



CARBON DIOXIDE IN AIR 



quantities, and that is the reason why lime-water, poured into 
a jar of ordinary air, does not at once become turbid. 

Pure country air contains only about 3 to 4 parts of CO2 per 
10,000 parts of air by volume. In towns where much coal is 
burnt, especially if there is no wind, this amount may rise to 
6 or 7 parts ; whilst in badly ventilated and crowded rooms, in 
which gas is also being burnt, the proportion may reach as much 
as 20 or 30 parts. 

Methods of Estimating the Amount of CO^ in Air. 
— The amount of CO^ in air can be found by aspirating a 
definite volume of dried air (not less than 20 litres) through (J 
tubes and bulbs filled with caustic potash, KOH, this absorbs 
CO2, and forms potassium carbonate, K^COo. 2 KOH + CO2 — 




Fig. 65. 

K2CO3 -f HgO. The tubes are weighed before and after the 
experiment, and the amount of carbonic acid absorbed is thus 
found. Fig. 65 shows the apparatus employed. A is an 
aspirator, consisting of two large glass bottles movable on a 
horizontal axis. The upper one is filled with water, w^iich, as 
it flows into the lower one, draws after it an equal volume of 
air which passes through the series of tubes and stopcock B. 
The first tube contains pumice stone, moistened wdth strong 
sulphuric acid, whilst the second is filled with dry calcium 
chloride ; these absorb all the aqueous vapour before the air 
passes through the next (J tube, in which most of the carbonic 
acid is absorbed, the air next bubbles through a strong solu- 
tion of caustic potash in the absorption bulbs, and passes, last 



LESSON XV CARBON DIOXIDE IN AIR 141 

of all, through another tube filled with small pieces of caustic 
potash, which ensures that no loss of water vapour takes place 
from the bulbs. 

Pettenkofer's Method of Estimating CO^ in Air. — 
This method is much more convenient and more easily carried 
out than the foregoing, and is generally used. A large cylin- 
drical stopped glass bottle, holding 10 litres or more, is used. 
This is filled with the air to be examined, and a measured 
volume of baryta-water of known strength is added. After 
shaking the liquid with the air, and allowing it to stand for 
half an hour, the whole of the CO^ is absorbed. 

Ba(OH). + CO, = BaCOo -f H,0. 

(Alkaline.)" " (NeutraL) 

Half the volume of liquid is now taken out by means of a 
pipette, and the amount of baryta remaining in solution is 
found by adding a solution of oxalic acid of standard strength 
until all the free baryta is neutralised. The baryta and oxalic 
acid solutions are made of such a strength that equal volumes 
of each will neutralise (see p. 83) each other, and so that 
I c.c. of baryta will absorb (say) i c.c. of CO2'. 

Example 

Volume of air taken for analysis 10 litres 

" baryta solution added (i c.c. solution = i c.c. CO2 

NTP) ^. 50 c.c. 

" oxalic acid solution required to neutralise 25 c.c. of 

the turbid baryta 22 c.c. 

" baryta solution combined with CO2 (25 - 22) X 2 = 6 c.c. 

" CO2 in 10,000 c.c. of air (10 litres) = 6 c.c. or 6 parts per 10,000. 

Water Vapour in the Air. — We have evidence of the 
presence of aqueous vapour in the air in the occurrence of 
dew, hoar frost, snow, and rain. A shallow dish full of water, 
if placed in the open air in summer, quickly evaporates into 
the air, and similar evaporation is constantly taking place from 
the sea, lakes, and rivers. If we bring a cold bright mirror 
into a warm room, the mirror becomes dimmed and bedewed 
with moisture. A person wearing spectacles in the open air 
in winter, finds, on coming into a warm gas-lit room, that his 
spectacles are dimmed so much as to render them temporarily 



142 



SNOW CRYSTALS 



useless, this is because the cold glass has condensed the water 
vapour present in the air of the room, and caused an artificial 
deposition of dew. 

The quantity of water which can be taken up by the air 
depends upon the temperature. At high temperatures air can 
take up a much larger amount of water as gas than at low 
temperatures. If warm moist air be cooled, it soon reaches a 
temperature at which it can no longer retain all its aqueous 
vapour as a gas, and this is deposited as liquid water. The air 

is then said to be saturated, 
and the temperature at which 
this occurs is called the dew- 
point. This deposition of liquid 
water takes place readily when 
minute particles of dust are 
present, and thus mist and fog 
are formed. Warm air ascend- 
ing from the sea, and heavily 
charged with water vapour, 
reaches a height at which the 
temperature is low enough to 
condense part of the water into 
clouds. If now these are blown 
into a much colder region, 
a further condensation takes 
If one cubic mile of air, saturated 
with water at 35° C. be cooled to 0°, it will deposit upwards 
of 140,000 tons of water as rain. If the condensation takes 
place in an atmosphere below the freezing point, snow or hail 
is produced. Snow is water vapour which has assumed the 
crystalline form in condensing in a cold atmosphere. Fig. 66 
shows the various forms of snow crystals. When condensa- 
tion takes place at the cold surface of the earth, on clear 
summer nights, dew is formed ; whilst, if the surface of the 
earth is below 0° C, a similar condensation takes place as hoar 
frost. One cubic metre of air can take up of water vapour, 




Fig. 66. 

place, and rain is produced. 



At qP C. 4.871 grams. 

" 5"C. 6.795 " 

" 10^ C. 9362 " 

" 15^ C. 12.746 "• 



At 20° C. 17.157 grams. 
" 30° C. 30.095 " 

40^ C. 50.700 
" 100-' C. 588.73 



LESSON XV ESTIMATION OF WATER VAPOUR 



143 



Estimation of Water Vapour in Air. — The increase 
in weight of the first two U tubes (Fig. 65), used in the esti- 
mation of CO2 in air, give also the amount of water present 
in the volume of air experimented with. A more simple 
method of ascertaining the amount of moisture in the air 
consists in the use of the wet and dry bulb hygrometer. 
It consists merely of two sensitive thermometers placed 
side by side (Fig. 67), the bulb of one is surrounded 
by a piece of muslin which is constantly kept moist by a 
strand of the muslin 
dipping into a little 
vessel full of distilled 
water. Evaporation of 
water is constantly tak- 
ing place from the moist 
muslin round the bulb. 
When water evaporates, 
heat is absorbed, and 
when the air is very dry 
the evaporation is rapid, 
and consequently the 
thermometer indicates 
a much lower tempera- 
ture than that of the air, 
as indicated by the dry- 
bulb thermometer. If 
the air be saturated with 
water vapour no further 
evaporation takes place, 
and both thermometers 
register the same temperature 




Fig. 67. 



From this " hygrometric 
break," or difference between the two thermometers, the amount 
of aqueous vapour present in the air can be computed from 
hygrometric tables. For other methods depending on the 
determination of the " dew-point," a work on physics must be 
consulted. 

Equilibrium of Chemical Composition of the Air — 
Action of Animals and Plants. — Our bodies are con- 
stantly undergoing a slow combustion ; the air which we expire 
is laden with carbonic acid gas and water vapour, owing 



144 ACTION OF ANIMALS AND PLANTS part II 

to the combustion or slow oxidation of the carbon and 
hydrogen of our bodies by the oxygen of the air. To prove 
this, we have only to breathe against a bright mirror, when it 
instantly becomes dimmed owing to the water vapour con- 
densing to minute drops of water. Similarly, if we breathe 
through a glass tube into a little lime-water, it soon becomes 
milky, showing the presence of carbonic acid gas. It may be 
asked, Why is it that all the fires burning, and all the animals 
breathing, do not use up the oxygen in the air, the amount of 
which, in fact, is not found to diminish? The answer is that 
the oxygen of the air is constantly being renewed, and the CO2 
diminished by the action of plants, the green parts of which 
contain chlorophyll^ which has the power of decomposing CO2 
in presence of sunlight, using the carbon * for the building up 
of its tissues, and liberating the oxygen for the use of man and 
animals. Thus we see that what is poison to man is the 
necessary food of plants, and whilst oxygen is a necessity of 
life for man, it is a mere bye-product of the life-processes of 
plants. Both processes, then, tend to produce an equilibrium 
in the chemical composition of the air. 

Organic Matter in the Air. — The motes which are seen 
dancing in the sunbeam are particles so small, that except when 
illuminated by the sunlight, they are not only invisible, but 
float in the air and are always more or less present in the 
atmosphere, except at great elevations. These motes contain 
minute particles of mineral matter, and also microscopic living 
organisms or their germs. These latter cannot, in strictness, 
be called chemical constituents of the air, and yet as they play 
a most important part in bringing about chemical changes, 
their existence and mode of action cannot be ignored even in 
elementary lessons on chemistry. These microbes or their 
germs cause such liquids as milk, beer, wine, etc., to turn sour 
on standing in the air. Exclude these microscopic organisms 
from the air, and these liquids remain sweet for any length 
of time, even when exposed to air, if thus purified ; give the 

* The tissues of plants contain large quantities of carbohydrates, or 
compounds which contain carbon combined with hydrogen and oxygen in 
the proportions in which they are combined in water, e.g. cellulose and 
starch, C1.2H20O10 ; cane sugar, CioHo^On ; grape sugar (glucose) and 
fruit sugar (levulose), C12H24O12. 



LESSON XV 



VENTILATION 



145 



organisms entrance, and the liquids soon turn sour. Fortunately 
most of the organisms which float in the air are not specially 
hurtful to animal life, but certain so-called ^' pathogenic " or 
'* disease '^ germs exist, and these are sometimes present in the 
air, and may give rise to disease. 

Ventilation. — The importance, for the due preservation 
of health, of breathing as pure air as possible, especially in- 
doors, cannot be too strongly insisted upon, and the chemist 
has found that when the amount of carbonic acid reaches 7 in 
10,000 of air, that air is not wholesome, not so much because 
this quantity of carbonic acid is hurtful, as because when this 
amount is reached in dwelling-rooms or w^ork-shops where 
human beings live or work, harmful exhalations from the body 
as well as microbic organisms are present, which exert a deteri- 
orating influence on the health. Hence it is necessary for the 
well-being and comfort of those inhabiting such rooms that a 
proper and continual renewal of the air should occur, and to 
eff'ect this without creating unpleasant draughts, is the object of 
a sound system of ventilation. 

If a number of people breathe the air, or if a candle flame be 
burning in a closed space where 
there is no means of ventilation, 
i.e. no outlet for impure air nor 
inlet for pure air, then that space 
becomes so laden with carbon 
dioxide that the people are suffo- 
cated, and the candle flame at last 
will not burn. 

ExPT. 64. — This experiment 
shows the need of ventilation in 
order to support the flame of a 
candle. A candle is placed in a 
glass dish which contains a layer 
of water, a narrow lamp chimney 
is now placed over the flame 
(Fig. 68), but as no air can enter 
the tube at the bottom, and as Ficr. 68. 

the products of combustion rise 

and prevent fresh air from entering at the top, the flame soon 
If now we relight the flame, and replace the lamp 




!»«! 



146 



VENTILATION 



PART II 




Fig. 69. 



glass, but this time furnished with a card-board partition, we 

shall find that the flame continues to burn, because the heated 

products of combustion 
can rise up one-half of 
the chimney and fresh 
air can enter down the 
other, as may be shown 
by holding a piece 
of smouldering brown 
paper near the inlet. 

One of the most im- 
portant means of ven- 
tilating dwelling-rooms 
is the chimney, wdiich 
usually contains a col- 
umn of heated air, w^hich 
ascends, being lighter 
than the air outside, 
whilst fresh air enters 
the room through every 

crevice or by means of the imperfect fitting of doors or windows, 

but preferably by properly constructed inlets. 
ExPT. 65. — Fig 69 shows an experiment 

which will illustrate this, one side of a box 

is removed and replaced by a sheet of glass, 

two holes are bored in the upper side of the 

box and these are covered wdth wide glass 

tubes. A lighted candle is placed under one 

of the tubes, and the heated air and products 

of combustion being lighter than the surround- 
ing air, ascend through it, whilst the colder 

and heavier air passes down the other tube to 

take its place, as may be seen by holding a 

piece of smouldering brown paper over the 

cold air inlet. 

It is in this manner that coal-mines are 

ventilated. The box in the figure may be com- 
pared to the workings of the mine, and the two 

tubes to the upcast and downcast shafts which ventilate the mine, 

a large fire being lighted at the bottom of the upcast shaft. 




Fig. 70. 



LESSON XV VENTILATION 147 

Thus a fire in a room is not only useful for its warmth, but 
also for its ventilating action. When gas-lights are burning, 
and people are breathing in a room, the hot and impure air 
ascends to the top of the room, as may easily be proved by 
standing on the table and breathing the upper layers of air. 
Hence, an outlet for these hot and impure gases should be 
provided by means of a suitable grating placed either in the 
outside wall, or leading into the chimney at points just below 
the ceiling. An inlet for fresh pure air should also be provided 
either by a grating in the outer wall below the level of the floor 
boards, or by raising the lower sash S of the window a few 
inches, and fitting a long piece of wood, B, in the lower aper- 
ture, this allows fresh air to enter the space between the two 
sashes as shown in Fig. 70, and purifies the air of the room 
without creating a draught, as the fresh air enters at a level at 
w^hich it will not be felt. 



What we have learnt 

In the fifteenth Lesson we have studied the various methods of prepar- 
ing nitrogen and examined its most characteristic properties. We have 
considered the most important facts regarding the atmosphere, viz., its 
weight and pressure, its extent and temperature. 

We have seen that the air is a mechanical mixture and not a chemical 
compound of nitrogen and oxygen, and have learnt the methods by which 
air is analysed, not only as regards its principal components, nitrogen and 
oxygen, but also as regards carbon dioxide and moisture. The effects of 
animal and vegetable life upon the atmosphere have been considered, and 
it was shown that each tends to support an equilibrium in the chemical 
composition of the air. Lastly, the principles of ventilation have been 
touched upon, and illustrated by various experiments. 



Exercises on Lesson XV 

1. Explain the various methods of preparing nitrogen. 

2. What is the action of chlorine on ammonia ? 

3. Describe several experiments to show the weight and pressure of the 
atmosphere. 

4. Write out an account of the various reasons for regarding air as a 
mixture and not a compound of nitrogen and oxygen. 



148 EXERCISES 



5. Describe the experiments, both volumetric and gravimetric, by which 
the composition of air has been ascertained. 

6. How is the amount of atmospheric carbon dioxide and water vapour 
ascertained gravimetrically ? 

7. Explain the effects of animal and plant life on the composition of 
the atmosphere. 

8. Describe experiments which illustrate the necessity for ventilating 
dwelling-rooms. 

9. Explain the action of the wet and dry-bulb thermometer. 

10. What is the proportional amount of CO2 existing in the open air; 
what is the proper limit in the air of rooms ? 

11. Explain the formation of dew, rain, hoar-frost, and snow. 

12. In what respect does air dissolved in water differ from atmospheric 
air ? 



LESSON XVI 

COMPOUNDS OF NITROGEN AND OXYGEN 

N2O, N2O2, N2O3, N2O,, and Np- 

NITROUS ACID, HNO2, NITRIC ACID, HNO3 

THE NITRITES AND NITRATES 

Nitric Acid, HNO3 

The starting-point for all the above compounds is nitre or 
saltpetre, KNO3, potassium nitrate, from which compound 
nitric acid, HNO3, is obtained. All the oxides of nitrogen may 
be prepared from nitric acid, and for this reason it will be 
treated first. 

Preparation. — Nitric acid is prepared by acting with 
sulphuric acid, H2S0^, on nitre, KNO3, or on Chili saltpetre, 
NaNOg, thus: — 

KNO3 + H,S0^ = KHSO, + HNO3. 

ExPT. 66. — In order to prepare nitric acid on a small 
scale, about equal parts of dry nitre crystals and concentrated 
sulphuric acid are placed in a stoppered retort and gently 
heated by a Bunsen burner (Fig. 71). The nitric acid distils 
over and is collected in a flask which is kept cool by a stream 
of cold water. The residue in the flask contains potassium 
hydrogen sulphate, KHSO^, or bi-sulphate of potash. On the 
large scale Chili saltpetre, NaNOo, a salt which is found as 
vast deposits in Chili, is used because it is cheaper. The 
salt is placed in a large iron cylinder, heated by a fire, and 

149 



ISO 



NITRIC ACID 



closed at the end by a circular stone flag, through a hole in 
which a large earthenware pipe is inserted for the purpose of 
carrying off the vapour of the acid to the condensers. 

Properties. — Nitric acid is a strongly fuming liquid, 
possessing a sharp acrid smell, and is colourless when pure, 
but is usually slightly tinged yellow, owing to the presence of 
oxides of nitrogen. It is a very corrosive substance, and 
hence it is sometimes called aqua fortis. When concentrated 
it burns the hands, and causes painful w^ounds, but if more dilute, 
it only stains the hands yellow. As its formula shows, it con- 




Fig. 71- 

tains 76 per cent of oxygen, and it acts as an energetic oxidising 
agent. This is seen by bringing a few copper turnings or a 
small piece of metallic tin into a test-tube containing a little 
of the acid. Red fumes are at once given off and the metals 
are oxidised, the tin to tin dioxide, SnOg, and the copper to 
copper nitrate, Cu(N03)2. In the same way, if strong nitric 
acid be poured upon sawdust, torrents of red fumes are emitted, 
and so much heat is given off that the sawdust is frequently 
inflamed. 

Tests for Nitric Acid. — In all cases, the salts of nitric 
acid (nitrates) must be mixed with concentrated sulphuric 



LESSON XVI THE NITRATES 151 

acid so as to liberate the nitric acid before the following tests 
can be applied. 

ExPT. 67. — (i) Copper turnings added to the acid, and 
the mixture warmed, give rise to dense red fumes possessing a 
characteristic smell, whilst the solution is turned blue owing 
to the formation of copper nitrate in solution. (2) A more 
delicate test consists in mixing the substance with concentrated 
sulphuric acid in a test-tube, cooling the mixture, and then 
pouring carefully dow^n the sides of the tube on to the surface 
of the mixture, a solution of ferrous sulphate (green vitriol) 
FeSO,^, w^hen a dark brownish-black ring forms at the surface 
where the liquids meet if a nitrate or nitric acid be present. 
This is known as the ring test. 

(3) Blue indigo solution is at once decolourised by nitric 
acid. 

Uses. — Nitric acid is largely used in the arts and manu- 
factures for dissolving silver, large quantities of silver nitrate 
being now used for photographic purposes. It is employed for 
making coal-tar colours, for preparing nitro-glycerine and gun- 
cotton, as well as collodion for photographic use. It is used 
also for etching on copper, and in the manufacture of sulphuric 
acid (seep. 207). 

The Nitrates 

The nitrates may be considered as nitric acid in which the 
hydrogen has been replaced by a metal, or group of elements 
equivalent to a metal, e.g. — 

HNO3 KNO3 (NHJNO3. 

Nitric acid (or Hydrogen nitrate). Potassium nitrate. Ammonium nitrate. 

Occurrence. — Although nitric acid itself does not occur in 
the free state in nature, its salts are widely distributed in the 
soil, and vast deposits of sodium nitrate, NaNOg, are to be 
found in Chili, whence its name Chili saltpetre. When nitro- 
genous animal matter is buried in the soil, the water draining 
from that soil is found to contain nitrates. The formation of 
these salts depends upon the action of certain minute 
organisms called bacteria, which have the power of converting 
the ammonia, given off during the decomposition of the 



152 THE NITRATES 



animal matter, into nitrous acid (HNO^), and nitric acid 
(HNO.) or their salts. 

Preparation. — In India and other hot countries, the 
nitrates which are thus produced from decomposing organic 
matter occur as an efflorescence on the surface of the soil, and 
it is the work of a certain caste of the natives (Sorawallahs, 
from sora^ nitre) to collect and purify the crude product by 
crystallisation. After recrystallisation, the salt is ready for 
sale. Potassium nitrate is now obtained on a large scale by 
the double decomposition of Chili saltpetre, with the naturally 
occurring potassium chloride (sylvine) found in large quantities 
at Stassfurt. 

NaNOo + KCl = KNO3 + NaCl. 

Chili saltpetre. Sylvine, Nitre. Common salt. 

The common salt is first deposited as crystals, whilst the 
mother-liquor contains the nitre wdiich crystallises out after- 
wards. 

Properties. — All nitrates are soluble in water, and when 
ignited usually decompose with the evolution of red fumes of 
the higher oxides of nitrogen, w^hilst oxides are left behind. 
Potassium nitrate, on heating strongly, is converted partially 
into the nitrite, KNO2, oxygen being evolved, whilst ammonium 
nitrate decomposes into nitrous oxide and water. 

Nitrate (NHJNO3 = N^ + 2 H,0. 
Nitrite (NHJNO2 = Ng + 2 Hp. 

As we have seen, the nitrite containing one atom of oxygen 
less in the molecule than the nitrate, evolves nitrogen under 
similar circumstances. 

ExPT. 68. — The oxidising action of nitre can be well shown 
by fusing some of the salt in a round bottomed hard glass 
flask, and dropping into it a small piece of charcoal, the 
carbon immediately bursts into flame, and continues to burn 
brilliantly until it is consumed by the oxygen in the nitre. 

Uses. — Potassium nitrate, KNO3, is chiefly used as an oxi- 
dising agent as it contains a very large proportion [47 per cent] 
of oxygen. It is therefore used in the manufacture of gun- 
powder and fireworks. It is also used in medicine and in the 



LESSON XVI NITROUS OXIDE 153 

pickling of meat. Chili saltpetre, NaNOo, is used in the 
manufacture of sulphuric acid (see p. 207), and as a manure, 
whilst silver nitrate, AgNOo, is used largely in photography, 
for the preparation of the chloride, bromide and iodide of 
silver. Lead nitrate, Pb(NO;5)o, ferric nitrate, Fe^(NOo),;, 
and aluminium nitrate, Al2(N0o),;, are used in dyeing and 
calico printing, whilst strontium nitrate, Sr(N0.3)2, and barium 
nitrate Ba(NO;5)2, are used for producing the crimson and green 
lights of pyrotechnic displays. 



THE OXIDES OF NITROGEN 

Nitrogen Monoxide or Nitrous Oxide 

Formula N.2O. Molecular Weight 43.7. Density 21.85 

Laughing Gas. — Nitrous oxide is sometimes called laughing 
gas, because when mixed with air and inhaled, it produces a 
peculiar and transient intoxicating eifect, and if inhaled in the 
pure state for a short time, it acts as an anaesthetic, that is, 
renders the person insensible to pain for a few minutes. It 
is, therefore, used for dental operations. This gas does not 
occur naturally. 

Preparation. — (i) We have already seen (under nitrogen) 
that when ammonium nitr//^ is decomposed by heat, nitrogen 
gas and steam are evolved. In the same way ammonium 
mirate^ prepared by neutralising nitric acid with ammonia, 
thus HN03 + NH3=(NH4)N03, a compound which contains 
one more atom of oxygen in the molecule, gives off nitrous 
oxide when decomposed by heat, thus : — 

(NH^)NO, == Ng + 2 H.p. (NHJNO3 = N,0 + 2 HgO. 

Nitrite. Nitrogen. Nitrate. Nitrous oxide. 

ExPT. 69. — Place about 20 grams of dry nitrate of am- 
monium in a flask furnished with a wide delivery tube, as used 
in the preparation of oxygen (Fig. 72). Heat the salt until it 
begins to decompose, and then regulate the flame of the lamp, 
otherwise the gas is evolved too rapidly. The pneumatic 
trough must be filled with warm water, because the gas 



154 



PROPERTIES OF NoO 



dissolves considerably in cold, but to a less extent in hot 
water. 

( 2) Nitrous oxide may also be obtained by the action of very 
dilute nitric acid on zinc, or other metals. 

Properties. — Nitrous oxide is a colourless gas, possessing 
a pleasant smell and sweet agreeable taste. It is condensed 
to a colourless liquid under a pressure of 32 atmospheres at 
0°. Poured into an open vessel the liquid is cooled down by 
its own evaporation to a temperature of — 100° C. Nitrous 
oxide has also been obtained in the solid state. Potassium and 
sodium take fire in this gas wdth formation of their peroxides, 



whilst the nitrogen is set free. 




Fig. 72. 

Experiments vrith Nitrous Oxide 

ExPT. 70. — Collect a test tube full of the gas, and insert 
a glowing chip, it will at once burst into flame, just as in 
the case of oxygen. 

ExPT. 71. — Having collected several jars of the gas, 
proceed as with oxygen, to burn phosphorus in a deflagrating 
spoon, the phosphorus will burn almost as brightly as in 
oxygen. 

ExPT. 72. — If a piece of sulphur, well alight, is intro- 
duced into the gas it continues to burn almost as brightly as 
in oxygen, but if only feebly burning, the flame is put out. 
The reason of this is that the gas must first be decomposed 



LESSON XVI DETERMIXATIOX OF COMPOSITION 155 

into its constituent elements, nitrogen and oxygen, before it 
can act as a supporter of combustion, and to effect this decom- 
position, a tolerably high temperature is necessary. 

ExPT. J2>' — Charcoal also burns as in oxygen. 

ExPT. 74. — Even a steel watch spring burns as in oxygen, 
if tipped with brightly-burning sulphur. 

Nitrous Oxide as a supporter of Combustion. — It is 
evident from the above experiments that nitrous oxide must 
first be decomposed into nitrogen and oxygen before bodies 
can burn in it, but we may ask. Why is it that this mixture of 
nitrogen and oxygen does not act as ordinary air. which is 
also a mixture of the same two gases? Let us inquire what 
would be the composition by volume of the mixture produced 
by the decomposition of nitrous oxide. 

Np = N. + O. 

2 vols. 2 vols. I vol. 

From the above equation we see that two volumes of the 
gas yield, on decomposition, two volumes of nitrogen and one 
volume of oxygen, so the mixture contains one-third or 33.3 
per cent of oxygen ; whereas atmospheric air only contains one- 
fifth or 20 per cent of oxygen, and it is to this larger proportion 
of oxygen that nitrous oxide owes its powers as a supporter of 
combustion. 

Distinguishing test for Nitrous Oxide and Oxygen. — 
The method by which these two gases may be distinguished 
from each other will be better understood after considering 
nitric oxide, and will be given under that gas, see p. 158. 

Determination of Composition of Nitrous Oxide. — 
(i) If a small piece of 
potassium is heated in a 
bent tube containing a 
measured volume of the 
gas, over mercury (Fig. 
73), the metal takes fire, 
combining with all the 
oxygen and liberating the ^^s- 73- 

nitrogen, and it is found that the volume of nitrogen left 
is the same after the experiment as the volume of gas 




156 



NITRIC OXIDE 



taken. Therefore nitrous oxide contains its own volume of 
nitrogen. 

(2) The density of nitrous oxide has been carefully deter- 
mined, and it is found to be nearly 22 times as heavy as hydro- 
gen, therefore two volumes weigh 44; but of this, 2 volumes or 
28 parts are nitrogen, hence the rest (44 — 28 = 16) is oxygen. 
Therefore we have 2 atoms (2 x 14) of nitrogen and one atom of 
oxygen (16) in the molecule, and the density (22) is half the 
moleculur weight (44). These experiments give the formula 
N^O to this gas. 



Nitrogen Dioxide or Nitric Oxide 



Formula* NO. Molecular Weight 29.8. Density 14.9 

Nitric oxide does not occur naturally. 

Preparation. — (i) Nitric oxide is best prepared by the 
action of strong nitric acid on metallic copper, the following 
equation expresses the reaction : — 

3 Cu + 8 HNO3 = 3 Cu(N03)2 + 2 NO + 4 Up. 

ExPT. 75. — Copper turnings are introduced into a flask so 
as to well cover the bottom, and these are covered with water ; 

strong nitric acid is then 
added through the thistle 
funnel (Fig. 74), when an 
evolution of gas begins. 
It will be noticed that the 
flask is first filled with a 
dark reddish-brown gas, 
but this colour soon dis- 
appears and a colourless 
gas is seen to be evolved, 
and collects in the s^as 




jars. 
The 



red colour first 



Fig- 74. 

noticed is due to the fact that nitric oxide has the power of com- 
bining with the oxygen of the air, which the flask first contains. 



* See determination of composition of nitric oxide, p. 158. 



LESSON XVI PROPERTIES OB' NITRIC OXIDE 157 



to form the higher oxides, N^O.^ and N^O^, which have a red 
colour, but the air in the flask is soon used up or expelled, and 
the gas is then seen to be colourless. 

ExPT. 76. — The gas thus obtained, is, however, not pure, 
but contains free nitrogen and nitrous oxide. To obtain it in 
the pure state the gas prepared as above must be led into a 
cold concentrated solution of ferrous sulphate, FeSO^ (green 
vitriol), with which it forms a peculiar compound of a deep 
blackish-brown colour. It is this compound which is produced 
in the ring test for nitrates. When this deep-coloured solution 
is heated the pure gas is given off. 

Properties. — Nitric oxide is a colourless gas which, on 
coming in contact with atmospheric air or oxygen, combines 
with the latter to form red fumes of the higher oxides, N^Og, 
and N^G^ (see Expts. ']^ and 78). It is not easily condensed 
to a liquid as it requires a pressure of 104 atmospheres at a 
temperature of -11° C. The gas is decomposed by heated 
metallic potassium which combines with the oxygen and 
liberates the nitrogen. 

Experiments with Nitric Oxide. — Having collected 
several jars of the gas prepared by the first method, we will 
make the following experiments. 

ExPT. ']^. — Remove the cover from one of the jars and 
expose the gas to the air. Deep brownish-red fumes (N^Og 
and N2O4), possessing a characteristic and disagreeable acrid 
suffocating smell, are noticed. 

ExPT. 78. — Pass oxygen into a second jar of the gas whilst 
still on the shelf of the pneumatic trough, deep red fumes are 
produced as .before, but these quickly disappear as they are 
very soluble in water. If the oxygen be added gradually to 
the pure gas, and in the proper proportion, all the gas will 
disappear owing to the red fumes dissolving in the water. 

ExPT. 79. — Phosphorus burns brilliantly in nitric oxide, 
but only when it is already brightly burning when brought 
into the gas. The flame of feebly-burning phosphorus, as well 
as those of sulphur and of a candle are, on the other hand, 
extinguished on plunging them into nitric oxide because the 
temperature of these flames is not sufficiently high to decompose 
this gas into its elementary constituents. 

ExPT. 80. — If a few drops of carbon disulphide, CS^,, be 



158 DETERMINATION OF COMPOSITION Part il 

poured into a cylinder full of the gas and shaken so as to 
allow the vapour to mix with it, it will be found that the 
mixture is very inflammable and burns, on applying a light, 
with a splendid blue and intensely luminous flame which is 
very rich in chemically active rays. 

Test for Nitric Oxide and distinguishing Test for 
Nitrous Oxide and Oxygen. — The above experiments 
ij^ and 78) will explain how free oxygen may be used as a test 
for nitric oxide, and this latter gas has been used not only as a 
test for free oxygen but also as a means of estimating its 
amount, and it was formerly used for the analysis of air. We 
now, however, possess better methods (see Air, p. 137). 
Nitrous oxide does not form red fumes with nitric oxide as 
oxygen does, they may thus be distinguished from each other 
although they both answer the glowing chip test for oxygen. 

Determination of Composition of Nitric Oxide. — 
The composition of nitric oxide is ascertained in the same 
way as that of nitrous oxide, viz. by heating a small piece 
of metallic potassium in a measured volume of the gas over 
mercury (Fig. 75). 

(i) It is found that the volume of nitrogen liberated 
is only half that of the nitric oxide taken. (2) The density 

of nitric oxide is found 
to be nearly 15, therefore 
2 volumes weigh 30, but 
this volume contains only 
I volume or 14 parts by 
weight of nitrogen, there- 
fore the oxygen is 30 — 
. 14 = 16 parts. This shows 

that the compound contains 
one atom of nitrogen combined with one atom of oxygen in the 
molecule, and its density (15) therefore is half its molecular 
weight (30). 

But why is the formula NO given to nitric oxide instead of 
NgOg which we have previously used for nitrogen ^/-oxide ? 

Because if we use N2O2 for the molecular formula then 
this gas will be an exception to the law (p. 55) that the 
density of any compound gas is half its molecidar weight. 
The density being 15, the molecular weight, according to this 




LESSON XVI N.2O3 AND NITROUS ACID 159 

law, is 30 and the formula is NO. It is as though the single 
molecule N^O^ had split up into two molecules of NO. The 
name nitrogen <^/-oxide is given to this gas because for the same 
weight of nitrogen it contains twice as much oxygen as nitrogen 
;;/^//-oxide. The remaining three oxides of nitrogen need only 
be considered here very shortly. 

Nitrogen Trioxide 

Formula N.2O3. Molecular Weight 75.5. Density 37.7 

This gas is formed together with the tetroxide when nitric 
oxide is exposed to the air or oxygen. It is obtained when a 
mixture of 4 volumes of nitric oxide and i volume of oxygen is 
allowed to pass through a hot tube. 

2 NO + O = N.Oo. 

Nitrogen trioxide is a dark-red gas and is easily condensed by 
a freezing mixture to a deep blue liquid at the ordinary pressure 
of the air, but even at — 2° the liquid trioxide is unstable, de- 
composing with liberation of nitric oxide. 

Nitrous Acid, HNO^. — Nitrogen trioxide is an acid-forming 
oxide, and comj bines with ice-cold w^ater to form nitrous acid. 

N,03 + H,0 = 2 HNO,. 

Nitrous acid is a very unstable substance, and has not been 
prepared in the pure state. Even its aqueous solution rapidly 
undergoes decomposition when heated, giving rise to nitric acid 
and nitric oxide, thus : — 

3 HNO. = HNO3 + 2 NO + H,0. 

The Nitrites or salts of nitrous acid are, on the other hand, 
very stable bodies, they are all soluble in water and give off red 
fumes of N2O0 when treated w^ith an acid. 

Test for Nitrites. — If a solution of a nitrite be added to a 
mixture of boiled starch with potassium iodide acidified with 
acetic acid, iodine is liberated and the starch is turned blue. 

HNO2 + HI = H,0 + NO + I. 

Nitrous acid. Hydriodic acid. Iodine. 



i6o NO2 AND N2O5 (NITRIC ANHYDRIDE) part il 

This reaction serves as a very delicate test for a nitrite, and is 
useful in testing water for nitrites, which may be taken as evi- 
dence of the previous contamination and gradual decomposition 
of animal matter. 



Nitrogen Tetroxide or Nitrogen Peroxide 

Formula NO^. Molecular Weight 45.7. Density 22.8 

Nitrogen peroxide has already been mentioned as being formed 
together with the trioxide. It may be prepared by mixing four 
volumes of nitric oxide with two volumes of oxygen and leading 
the red fumes into a tube surrounded by a freezing mixture. 

2NO + 02=2N02. 

It may also be formed by the decomposition of lead nitrate by 
heat. 

Pb(N0o)2 = PbO + 2 NO2 + O. 

It may also be prepared by heating nitric acid with arsenic 
trioxide, AS2O3, a mixture of nitrogen trioxide and peroxide is 
evolved which is condensed in a tube surrounded by a freezing 
mixture ; by passing oxygen through this liquid the trioxide is 
oxidised to tetroxide. Nitrogen tetroxide solidifies to colourless 
crystals at — 9°, slightly above this temperature the liquid also 
is colourless, but at 10° it attains a decided yellow colour which 
darkens further at higher temperatures. 

Nitrogen Pentoxide or Nitric Anhydride (N^O-). — This 
compound is a white crystalline solid obtained by removing the 
elements of water from nitric acid by means of phosphorus 
pentoxide, P20^, a substance which has a great power of ab- 
stracting water. 

2HN03 = N205 + H2O. 

Nitrogen pentoxide is an unstable body, and is not used in the 
arts, it is interesting as being the highest oxide of nitrogen, and 
the aiihydride of nitric acid, uniting with great energy with 
water to form this acid. 



LESSON XVI SUMMARY AND EXERCISES i6i 



What we have learnt 

In the sixteenth Lesson we have learnt the method of preparing nitric 
acid, nitrates, and the five oxides of nitrogen, and the mode of showing 
their properties and composition experimentally. 

Nitric acid is prepared from nitre and sulphuric acid ; nitrous oxide 
from ammonium nitrate, and nitric oxide by the action of nitric acid on 
copper. 

Exercises on Lesson XVI 

1. How is nitric acid prepared ? How would you proceed to show its 
properties ? 

2. How is nitre prepared from Chili saltpetre ? 

3. What is the result of strongly heating (a) ammonium nitrite, (^) 
ammonium nitrate ? 

4. Why is nitrous oxide a good supporter of combustion ? How would 
you distinguish it from oxygen ? 

5. I want 12 litres of nitric oxide at 14- C. and 738 mm. How much 
copper and nitric acid must I take ? 

6. How can you prove that the formulas N2O and NO represent the 
composition of nitrous and nitric oxide ? Why is NO called nitrogen 
dioxide ? 

7. How would you test a sample of water for nitrites ? 

8. What volume of nitrous oxide at 15^ C. and 770 mm. can I get by 
decomposing 400 grams of ammonium nitrate, and what volume of oxygen 
at NTP does the gas contain ? 

9. How is nitrogen pentoxide prepared ? Give an equation showing 
its action upon water. 

10. What is laughing gas ? How is it prepared ? 

11. Why is the formula NO2 given to nitrogen tetroxide ? 

12. A w^hite crystal is said to be nitre ; how would you ascertain whether 
this is so or not ? 

13. \Vhat weight of sylvine is required to decompose completely one ton 
of Chili saltpetre, and how much nitre will be theoretically formed ? 



LESSON XVII 

COMPOUNDS OF NITROGEN AND HYDROGEN, 

AMMONIA (NHg), AND THE AMMONIUM (NHJ, 

COMPOUNDS 

Nitrogen combines with hydrogen to form three compounds : 
(i) Ammonia, NHg. (2) Hydrazine or diamide, N^H^. (3) 
Azoimide, NgH. Of these the first is the most important and 
is the only one which we shall here study. 

Ammonia 

Formula NH3. Molecular Weight 16.9. Density 845 

We have learnt that the animal body contains substances 
into which nitrogen enters as a component part ; certain por- 
tions of plants, especially the fruit, seeds, and juice also contain 
nitrogenous compounds. Now, whenever these nitrogenous 
materials, whether of plants or of animals, are decomposed 
either naturally in the air or by heating them, ammonia is 
given off. The peculiar pungent smell of this compound is 
noticed if we heat a bit of cheese in a test-tube, and its 
presence can be further shown by thrusting a piece of moist- 
ened red litmus paper into the test-tube when the red colour 
will be changed to blue, for ammonia is a gas which has an 
alkaline reaction. This same smell is often noticed in urinals 
and stables, showing that animal nitrogenous matter has there 
been converted into ammonia. A common name for ammonia 
is spirits of hartshorn, this substance having, in former days, 
been prepared by heating horn. Again, the name ammonia is 

162 



LESSON XVII 



PREPARATION OF AMMONIA 



163 



derived from the temple of Jupiter Ammon in the Libyan 
desert, because it was there that the Arabs first prepared a salt 
of ammonia — Sal-ammoniac — by heating dry camePs dung. 

Ammonia is now almost exclusively obtained as a bye- 
product in the destructive distillation of coal in making coal- 
gas, w^here it collects as ammoniacal liquor. This liquor is 
neutralised by hydrochloric acid, and on evaporation, a solid 
salt is obtained known as Sal-ammoniac, NH^Cl (a compound 
of ammonia, NHo, and hydrochloric acid, HCl), and from this 
compound ammonia gas can be readily prepared. 

Preparation. — Ammonia gas is obtained in the laboratory 
by heating sal-ammoniac with lime. 

ExPT. 81. — A mixture is made of two parts of powdered 
quicklime and one part of sal-ammoniac. This is placed in a 
flask and covered with a layer of dry quicklime, CaO, which serves 
to absorb the moisture given off during the decomposition. 



CaO + 2(NHJC1 

Lime. Sal-ammoniac. 



= CaClg + 

Calcium Chloride. Ammonia. 



2NH3 + 



up. 



The mixture is heated, and 
the gas led off by a delivery 
tube as in Fig. 76. 

Ammonia gas cannot be 
collected over water be- 
cause of its great solubility ; 
with water it forms spirits 
of hartshorn or liquor 
ammonia of the shops ; but 
it can be collected over 
mercury, or, more simply, 
by upward displacement 
(see Fig. 76) . 

ExPT. 82. — To ascer- 
tain when the jar is filled 
with ammonia it is only 
necessary to hold a stopper 
or glass rod moistened with 
strong hydrochloric acid 
below the mouth of the ^^' ^ ' 

bottle, and if the jar is full, and ammonia gas is escaping 




164 



PROPERTIES OF AMMONIA 



dense white fumes of sal-ammoniac will be seen to form round 
the stopper (NH3 + HCl = NHp). 

Properties. — Ammonia is a colourless gas possessing a 
peculiar pungent odour well know^n as that of smelling salts, 
w^hich contain it. When the pure gas is inhaled or the strong 
liquid swallowed death often ensues. It is lighter than air, 
its density being 0.59 (air=i). It is exceedingly soluble 
in water, one volume of water absorbing 1148 vols, of the 
gas at 0°, whilst at 20° 741 vols, of the gas are absorbed 
under a pressure of 760 mm. 

ExPT. 83. — If a jar of the gas is brought under the sur- 
face of water contained in the pneumatic trough ; on removing 
the cover the water is seen to rush up 
and fill or nearly fill the jar, according 
as the whole or only a part of the atmo- 
spheric air has been expelled by the 
ammonia gas. 

ExPT. 84. — A better way of showing 
this is to fill a round bottomed flask with 
the gas by displacement, and insert a 
stopper, through which passes a piece 
of glass tubing, ending in a jet passing 
up into the middle of the flask, and 
outwards for about a foot so as to dip 
into a beaker of water (Fig. 'j^^. On 
cooling the gas by dropping a small 
quantity of ether on the outside of the 
flask, a little water enters. This dis- 
that the pressure inside the flask is 




Fig. 77. 



solves so much gas 
greatly reduced, and the water rushes into it in the form of 
a fountain. If the w^ater in the beaker be first coloured 
red wath litmus, the alkaline character of the solution may 
be demonstrated by the change of colour to blue as the liquid 
enters the flask. 

ExPT. 85. — If the delivery tube be allowed to dip into a 
bottle half filled with distilled water, it will be noticed that 
although the gas may be bubbling through the wash-bottle 
very rapidly, yet no gas passes through the water (Fig. 78), as 
all the bubbles are absorbed almost as soon as they reach the 
liquid. In this manner the aqueous solution is prepared. The 



LESSON XVII AMMONIA FREEZING MACHINE 



165 




Fig. 78. 



saturated solution of ammonia is lighter than water, and has a 
specitic gravity of about 0.880. 

Ammonia gas can be liquefied by exposure either to cold 
(about — 34^), or to pressure (about 7 atmospheres at the usual 
temperature of the air), and the liquid freezes 
to an ice-like solid if further cooled below 
— 75^. Ammonia does not support combus- 
tion under ordinary circumstances, but if 
mixed with oxygen and a light applied to 
the end of the tube from which the mixed 
gases escape, the hydrogen of the ammonia 
unites with oxygen and the mixture burns 
with a bright flame tinged yellow by the 
presence of the nitrogen, which, however, 
does not burn. 

Ammonia Freezing Machine. — We 
have seen on p. 108 that when steam is condensed to water the 
large amount of heat necessary for the existence of the gaseous 
condition becomes free, and vice versa, that w^hen water is con- 
verted into steam heat is absorbed (see p. 109). The same 
thing happens (to a greater or less extent) whenever a gas 

becomes a liquid or a liquid 
a gas ; and this can be made 
use of in the case of ammonia 
for the artificial production 
of ice. All that is needed 
is to have two strong iron 
vessels connected together 
with a pipe ; into one of 
these (the retort) is placed 
an aqueous solution of am- 
monia saturated at o^, and 
the whole made air-tight. 
If we wish to make ice, the 
water to be frozen is placed 
^^^- 79* in the inside of the second 

vessel (the receiver) which is hollow, and this vessel plunged 
into a bucket of cold water. Heat is now applied to the 
retort, the effect of this is to drive out the ammonia gas 
from the solution, this passes over into the receiver and grad- 




i66 EXPERIMENTAL DETERMINATION OF part ii 

ually collects there until the pressure of the accumulated gas 
becomes so great (about lo atmospheres) as to cause lique- 
faction of the ammonia. The liquid ammonia then collects in 
the hollow receiver. As soon as the aqueous solution in 
the retort has become hot and has given off the greater 
part of its ammonia, the position of the vessels is altered ; 
the retort is placed in cold water, and the receiver in the 
air, but surrounded with flannel. What takes place now? A 
reabsorption of ammonia in the cooled water begins, and a 
consequent evaporation of the liquid ammonia in the receiver, 
so much heat being absorbed in this evaporation that the water 
placed in the inside of the receiver is frozen. This process can 
be made continuous on a large scale, and thus tons of water can 
be frozen at a comparatively small cost. It will be seen that 
thus by burning coal we make ice. 

ExPT. 86. Composition of Ammonia. — If we pass a 
series of electric sparks through a measured volume, say 40 
c.c. of dry ammonia gas collected in a eudiometer (see p. 93) 
over mercury, the volume of the gas is seen to become greater, 
until at last no further increase takes place. On measuring 
the gas we now find that the volume is exactly double what 
it was to begin with, i.e. 80 c.c. This is because the ammonia 
has been decomposed into its elements, viz. nitrogen and 
hydrogen. 

2NH3 = N^ + 3H2. 

4. vols. 2 vols. 6 vols. 

If we now add a measured quantity of oxygen sufficient to 
combine with all the hydrogen, to form water when a spark 
is passed into the mixed gases, the diminution in bulk which 
follows the passage of the spark is due to the formation of 
liquid water, and for this, as we know, 2 volumes of hydrogen 
unites with i of oxygen (p. 94). We find in our experiment 
that this loss of volume amounts to 90 c.c, of which | or 
60 c.c. must be hydrogen. Therefore there must have been 
80 — 60 = 20 volumes of nitrogen. Hence 40 volumes of 
ammonia contain 20 volumes of nitrogen and 60 of hydrogen, 
or 2 volumes contain i volume of nitrogen and 3 volumes of 
hydrogen, and NHo represents the composition of ammonia. 
ExPT. 87. — Another method of showing the composition 



LESSON XVII 



COMPOSITION OF AMMONIA 



167 




of ammonia consists in filling with chlorine a long glass tube 
closed at one end, and stoppered at the other, the stoppered end 
communicating with a smaller stoppered tube (Fig. 80) . The 
stopper dividing the tube into two portions is not pierced through, 
but contains a cavity capable of holding a few drops of liquid. 

When the long tube is completely hlled with chlorine, it is 
divided into three equal portions by elastic bands fitting tightly 
on the tube, and a strong solution of ammonia is placed in the 
upper short portion, and the 
stopcock is turned so as to 
admit a few drops of it to 
the lower tube. The en- 
trance of the ammonia is 
accompanied by a small 
lambent yellowish green 
flame at the point where 
the drop enters the gas. This is due to a 
combination between the hydrogen of the 
ammonia with the chlorine to form hydro- 
chloric acid, nitrogen being liberated. When 
ammonia has been added drop by drop until 
all the chlorine has disappeared, a little sul- 
phuric acid is introduced in the same way as 
the ammonia was, so as to absorb the excess 
of ammonia gas. We know, however (p. 181), 
that chlorine combines with its own volume 
of hydrogen to form hydrochloric acid, and 
therefore we must have used the same volume 
of hydrogen as we took of chlorine to com- 
mence with. It only now remains to find 
out what is the volume of nitrogen remaining, 
and this is easily done by inverting the tube 
in water and taking out the stopper ; when water enters the tube 
to fill the partial vacuum and on equalising the pressure inside 
and outside the tube, it is seen that the nitrogen remaining 
occupies just one-third the length of the tube, showdng that i 
volume of nitrogen had been combined with 3 volumes of 
hydrogen. 



Fig. { 



2NH3 + 3CU 
4 vols. 6 vols. 



N, + 6HC1. 

2 vols. 



i68 THE AMMONIUM SALTS part ii 

The Ammonium Salts. — Ammonia unites with acids to 
form salts, called the salts of a7/i7no7imfn. .Because these salts 
so closely resemble the salts of the alkali metals, potassium 
and sodium, ammonia is termed the volatile alkali. The 
analogy in constitution of the salts of ammonia with those of 
potash and soda is exhibited by supposing that the former 
contain a component NH^ (ammonium), w^hich acts like the 
metals K or Na. Thus, if we write sal-ammoniac NHo.HCl, 
the analogy w^hich exists between it and potassium chloride, 
KCl, is not evident ; whilst if we represent the first salt by the 
formula (NH4)C1, in which NH^ displaces K, the analogy is 
at once seen. So we have — 

Salts of Potassium. Salts of Ammonium. 

Potassium chloride, KCL Ammonium chloride (NH4)C1 

Potassium sulphate, K2SO4 Ammonium sulphate (NH4)2S04 

Potassium nitrate, KNO3 Ammonium nitrate (NH4)N03 

The quasi-metal — that is, a body acting like a metal, but not 
really one — has not been obtained in the free state. 

What we have learnt 

In the seventeenth Lesson we learnt how to prepare ammonia by heat- 
ing sal-ammoniac with lime and collecting the gas by upward displacement, 
as it is lighter than air. We have seen that the principal properties of 
ammonia are its solubihty in water, its alkaline properties enabling it to 
neutralise acids to form ammonium salts, which are analogous to the cor- 
responding potassium salts, and contain the group of atoms (NH4) which 
acts like a metal, and is called ammon/«;;z. 

We have learnt two methods of demonstrating the composition of 
ammonia, viz. (i) by decomposing the gas by electric sparks in a eudio- 
meter, and then finding the amount of nitrogen and hydrogen produced, 
by exploding the mixture with an excess of oxygen ; (2) by gradually add- 
ing strong ammonia solution to a measured volume of chlorine, which 
combines with the hydrogen bulk for bulk and liberates a volume of 
nitrogen equal to one-third the bulk of the chlorine. The ammonia freez- 
ing machine has also been described, and its mode of action explained. 

Exercises on Lesson XVII 

1. How would you prepare ammonia gas, and demonstrate its most 
characteristic properties ? 

2. What is the action of ammonia on (i) hydrochloric, and (2) sul- 
phuric acid, (3) chlorine, (4) water ? Give equations. 



LESSON XVII SUMMARY AND EXERCISES 169 

3. I want 10 litres of ammonia gas at NTP, what weight of ammonium 
chloride must I use ? 

4. Describe in detail two methods by which the composition of ammonia 
can be ascertained. 

5. Describe and explain the action of the ammonia freezing machine. 

6. Write down the formulas of several ammonium compounds, and 
show their analogy to the potassium compounds. 

7. How would you demonstrate the great solubility of ammonia gas in 
water ? 

8. From what source is ammonia derived on a large scale ? 

9. What is the weight of 746 c.c. of ammonia gas measured at 10° C. 
and 760 mm. ? 

10. What is the volume of ammonia (NTP) which, when passed intc 
4 grams of pure H2SO4 of 50 per cent strength, will neutralise the acid ? 

11. What volume of nitrogen at 14° R. and 560 mm. will be left after 
acting with ammonia on 486 c.c. of chlorine at 5° C. and 780 mm. ? 



LESSON XVIII 

CHLORINE, HYDROCHLORIC ACID, AND THE 
CHLORIDES 

Chlorine 

Symbol CI. Atomic Weight 35.2. Density 35.2 

Chlorine gas was first obtained and its properties examined 
by Scheele in 1774; he prepared it by the action of hydrochloric 
acid, HCl, on manganese ore, containing the dioxide, Mn02. 

Sir Humphry Davy in 18 10 first satisfactorily proved the ele- 
mentary nature of the gas and gave it its present name from its 
colour, X'^wpo? greenish-yellow. 

Occurrence. — Chlorine does not occur in the free state in 
nature, but its compounds are widely diffused ; thus common 
salt, which is present in such large quantities in sea-water, is 
sodium chloride, NaCl ; rock salt also has the same composition. 
Potassium chloride, KCl, or sylvine, is found in large quantities 
as a natural deposit at Stassfurt in Germany. Other chlorides 
will be mentioned under ''^The chlorides." 

Preparation. — (i) Chlorine gas is easily prepared by the 
action of strong hydrochloric acid, HCl, on black oxide of man- 
ganese, manganese dioxide, Mn02, manganese chloride being 
also produced. 

MnO^ 4- 4 HCl = MnCl^ + 2 H.O + Cl^. 

This reaction depends on the formation, in the first instance, 
of manganese tetrachloride, MnCl^, for when manganese dioxide 
and cold concentrated hydrochloric acid are brought in contact, 

170 



LESSON XVIII PREPARATION OF CHLORINE 171 



a dark-brown solution is formed which, on heating, decomposes 
into manganous chloride and chlorine, thus : — 

MnO, + 4 HCl = MnCl^ + 2 H^ 
MnCl^ = MnCl, + Cl^. 




Fig. 81. 

This process is employed for the manufacture of chlorine on a 
large scale for making bleaching powder. 

ExPT. 88. — The apparatus used for the preparation of chlo- 
rine in the laboratory, is shown in Fig. 81. The manganese 
dioxide in small lumps is first introduced into the large flask and 
well covered with the strong acid. The flask is connected by a 



172 PREPARATION OF CHLORINE part ii 

delivery tube bent twice at right angles with a wash-bottle (d) 
half filled with water, to absorb the vapour of. hydrochloric acid 
which is carried over with the chlorine ; it next passes through 
strong sulphuric acid in wash-bottle a, and along the inclined 
tube containing pumice stone moistened with strong sulphuric 
acid, which serves to deprive the gas of aqueous vapour. The 
gas is then collected in the gas jar e, by downw^ard displacement, 
as It is about two and a half times as heavy as air, and is soluble 
in water. The tube <^, which dips under caustic soda solution, 
acts as a safety valve in case the evolution of gas becomes too 
rapid. It is easily seen when the gas jar is full of the gas, by its 
green colour. As chlorine has a very irritating and poisonous 
effect on the nose, mouth, and throat, ^/iis experiment and all 
others with chlorine^ should be inade in a " draught chai7iber^^ as 
shown in Fig. 8i, so arranged that the apparatus can be enclosed 
by a sliding glass door, and all fumes carried away to a flue lead- 
ing to the chimney. 

(2) It is sometimes convenient to generate the hydrochloric 
acid from common salt, NaCl, and strong sulphuric acid, H9SO4, 
in intimate contact with the manganese dioxide. In this 
reaction the hydrochloric acid first formed is at once decom- 
posed by the manganese dioxide with evolution of chlorine, 
thus : — 

4 NaCl + 3 H2SO4+ Mn0.2= Na.2S04 + 2 NaHS04+ MnCl2+ CI2+ 2 H.2O. 

ExPT. 89. — In this method 11 parts by weight of common 
salt afe mixed with 5 parts of manganese dioxide and 14 parts 
of sulphuric acid, diluted with an equal bulk of water, are added. 
On gently heating the mixture, chlorine is given off regularly. 

(3) Chlorine may also be prepared by the electrolysis of 
hydrochloric acid. 

2 HCl = H2 -f CI2. 

(See determination of composition of hydrochloric acid, p. 179.) 

(4) Chlorine gas is evolved when an acid such as hydro- 
chloric or sulphuric is added to bleaching powder, CaOClg 
(see p. 185), thus : — 

CaOClg + 2 HCl = CaCl, + H,0 + Clg, 
CaOClg + H2SO4 = CaSO^ + H.G + CI,. 



LESSON XVIII PROPERTIES OF CHLORINE 173 

(5) When a mixture of air and hydrochloric acid gas (HCl) 
is strongly heated, the hydrogen of the acid is oxidised to 
water by the oxygen of the air, whilst chlorine is liberated, 
thus : — 

2 HCl + O = H^O + Cl^. 

This reaction is used on the large scale for the economic pro- 
duction of chlorine, and is known as Deacon's process. It is 
employed in the manufacture of bleaching powder (see p. 185). 

Properties. — Chlorine is a transparent gas of a greenish- 
yellow colour, and possessing a most disagreeable and power- 
fully suffocating smell, which when the gas is present in small 
quantity only, resembles that of seaweed, but in large quantities 
produces violent irritation in the nose, mouth, and throat, 
giving rise to coughing and inflammation of the mucous mem- 
branes. If inhaled in the pure state it causes death, i litre 
of chlorine weighs 0.0899 ^ 35-2 = 3-1645 grms. Under a 
pressure of six atmospheres at 0°, or when exposed to a 
temperature of — 34^ at the ordinary atmospheric pressure, 
chlorine is condensed to a yellow liquid. At — 102° the liquid 
solidifies to a yellow crystalline mass. 

Chlorine is soluble in water to which it imparts its peculiar 
properties, one volume of water dissolves two volumes of 
chlorine, the solution is known as chlorine water and is used 
in the laboratory for various purposes. As chlorine also com- 
bines with mercury, it cannot be collected either over that metal 
or over water, but may easily be obtained by downward dis- 
placement (Fig. 81), as it is two and a half times as heavy 
as air. 

Combustions in Chlorine 

ExPT. 90. — Chlorine combines energetically vvdth hydrogen, 
and if we lower a jet of burning hydrogen into a jar of chlorine, 
the flame continues to burn, but instead of gaseous water being 
the product of combustion, dense fumes of hydrochloric acid 
are produced. The presence of hydrochloric acid may be 
shown by holding a piece of moistened blue litmus paper 
in the fumes which are given off. It is immediately turned 
red. 



174 COMBUSTIONS OF CHLORINE part ii 

The following experiments show the great power chlorine 
possesses, not only of combining with free hydrogen but with 
hydrogen in a state of combination. 

ExPT. 91. — When a piece of filter paper moistened with 
turpentine, CiqH^^, is plunged into chlorine gas, so much heat 
is evolved in the combination of the chlorine with the hydrogen, 
that some of the turpentine takes fire, clouds of soot and fumes 
of hydrochloric acid being evolved. 

ExPT. 92. — Another experiment is to plunge a lighted taper 
into a jar of chlorine, the taper continues to burn but with a 
smoky dull red flame which emits dense fumes of hydrochloric 
acid as well as a cloud of black soot. This is because the 
chlorine combines with the hydrogen of the wax and liberates 
the carbon. 

ExPT. 93. — Moist chlorine also combines directly with metals, 
sometimes with evolution of light and heat. Thus if a few 
leaves of Dutch metal (copper in thin leaves) are placed in a 
round bottomed flask furnished with a glass stopcock, and the 
flask be then exhausted of air, it will be found on opening the 
stopcock and admitting chlorine that the copper leaf will take 
fire and burn with the evolution of dense yellow fumes of copper 
chloride (CUCI2). 

ExPT. 94. — Similarly, if finely powdered antimony be thrown 
into the gas a shower of sparks accompanies the formation of 
antimony chloride (SbClg). 

ExPT. 95. — If sodium be melted in a deflagrating spoon 
and plunged into the moist gas, it takes fire, burning brightly 
with the formation of sodium chloride (NaCl) . 

ExPT. 96. — A small piece of phosphorus placed in a defla- 
grating spoon, and plunged into chlorine, first melts and soon 
bursts into flame, with formation of the chlorides of phosphorus 
(PClgandPCl,). 

Chlorine also abstracts the hydrogen from sulphuretted hydro- 
gen, H2S, olefiant gas, C2H4, water, H2O, and other hydrogen 
compounds. 

H2S 4- CI2 = 2 HCl + S. 
C2H, + CI4 = 4 HCl + Cg. 
H2O + CI2 = 2 HCl + O. 



LESSON XVIII BLEACHING BY CHLORINE 175 

Bleaching Action of Chlorine 

We have seen above that chlorine has the power of decom- 
posing water with liberation of oxygen. Chlorine water when 
exposed to sunlight soon loses its yellow colour, and hydro- 
chloric acid is formed in the solution. At the moment of 
liberation of the oxygen, that gas has remarkably active 
bleaching powers ; thus if a piece of turkey red calico be dipped 
in chlorine water, the part immersed is soon decolourised. The 
oxygen which thus oxidises the colouring matter, is said at the 
moment of its liberation to be in the nascent state, and this 
difference is probably due to its being in its free atojnic 
condition, whilst in ordinary oxygen the free atoms have united 
to form the less active molecules. 

(Nascent oxygen) 2 O = O2 (Ordinary free oxygen) . 

This view is confirmed by the fact that dry chlorine does not 
possess bleaching power, and a piece of turkey red cloth may 
be kept in the dry gas without losing its colour, whilst if 
moisture be afterwards admitted the colour quickly disappears. 
Test for Chlorine. — Chlorine may be recognised by its 
characteristic smell, and by the fact that it colours iodised 
starch paper blue owing to the liberation of iodine. 

KI + C1 = KC1 + I. 

It may also be recognised by its bleaching action and by its 
characteristic yellow colour. With silver nitrate, chlorine and 
all the soluble chlorides produce a white precipitate of silver 
chloride, which is insoluble in nitric acid but soluble in ammonia. 

AgNOg + NaCl = AgCl +NaN03. 



CHLORINE AND HYDROGEN 
Hydrochloric Acid (Muriatic Acid) 

Formula HCl. Molecular Weight 36.2. Density 18.1 

Only one compound of chlorine and hydrogen is known, 
namely HCl. Glauber originally obtained this acid by the 



176 



HYDROCHLORIC ACID 



action of sulphuric acid on common salt about the year 1648. 
Priestley first obtained the gas itself by collecting it over 
mercury in the pneumatic trough. Before that time it was 
only known in solution, it being excessively soluble in water. 
But it was not until 18 10 that Sir Humphry Davy proved it 
to be composed of hydrogen and chlorine only. 

Occurrence. — The gas is not found free in nature except 
in small quantities in the gases evolved from certain active 
volcanoes. The occurrence of its salts has already been 
mentioned under chlorine (see also The Chlorides, p. 184). 

Preparation. — (i) Hydrochloric acid is usually prepared 
by the action of concentrated sulphuric acid on common salt. 
The reaction takes place in two stages, in the first of which 
one molecule of each substance reacts on the other, thus : — 

(i) NaCl + H,SO^ = NaHSO, + HCl, 
(2) NaHSO^ + NaCl = Na^SO^ + HCl. 

In the second reaction, the sodium hydrogen sulphate, or 
bisulphate of soda decomposes another molecule of salt with 




Fig. 82. 

formation of another molecule of hydrochloric acid. For the 
second reaction a higher temperature is required. 

ExPT. 97. — The apparatus used is exactly like that used 
for chlorine, and the method of preparation the same, hydro- 
chloric acid gas being collected by downward displacement. 

(2) Hydrochloric acid may also be produced by the direct 
combination of its elements, but this is not a convenient way 
of preparing it. To show the combination, a small jar of 
hydrogen is inverted over a similar jar of chlorine in the dark 
after allowing the gases to mix, a lighted taper brought to the 
mouth of the jar will cause their instant combination, which is 



LESSON XVIII 



MANUFACTURE OF HCl 



177 



accompanied by an explosion (see also explosive combination of 
chlorine and hydrogen, p. 181). 

Manufacture of Hydrochloric Acid. — Hydrochloric acid 
is prepared on a very large scale as one of the bye-products 
in the manufacture of soda-ash (alkali) or carbonate of soda. 
In the alkali works 10 cwt. of salt is placed in a large iron pan 
a (Fig. 82), heated by a fire b placed underneath, and covered by 




a dome or arch of brickwork. Upon this salt about an equal 
weight of sulphuric acid (sp. gr. 1.7) is run. Hydrochloric 
acid gas is then rapidly given off and led from the furnace 
by means of glazed stoneware pipes to high towers (show^n 
in Fig. 83) built of bricks, or Yorkshire flag, soaked in tar 
and clamped together and filled with bricks or coke. Down 
the interior of the tower a stream of water is allowed to 
trickle from a cistern (F) placed on the top. The hot gas 
from the furnace is drawn off by the draught of the chimney 



178 AQUA REGIA part ii 

(E), and passes at (A) into the tower (B), where the greater 
part is absorbed ; then the undissolved gas passes down the pipe 
(C) into the '^ exhaust '^ tower where, in ascending, it again 
meets with a stream of water, and thus every trace of hydro- 
chloric acid gas is absorbed. Formerly, the escape of this 
acid gas from alkali w^orks was a great source of nuisance and 
damage to neighbouring property. Now, under careful inspection, 
the condensation of the acid is so perfect that no perceptible 
injury is effected, whilst the manufacturers save a valuable pro- 
duct which was at one time allowed to escape into the air. 

In the pan a the first reaction takes place, in which the acid 
salt is formed, thus : — 

NaCl + H^SO^ = NaHSO^ + HCl. 

When this has taken place, the mass is raked on to the hearths 
of the reverberatory furnaces dd, and heated more strongly by 
the fires cc, when the further reaction is completed, viz., 

NaHSO^ + NaCl = Na^SO^ + HCl, 

and thus salt-cake or sodium sulphate, Na^SO^ is produced, and 
the whole of the hydrochloric acid given off. 

Properties. — Hydrochloric acid is a colourless gas possessing 
an acrid choking smell. It can be condensed to a colourless 
liquid at —4, under a pressure of 25 atmospheres. The gas fumes 
strongly in the air combining with the aqueous vapour there 
present. It is excessively soluble in water, one volume of water 
at o^ dissolving 503 vols, of the gas at NTP. The gas is heavier 
than air, and may be collected by downward displacement. Like 
all acids it turns blue litmus red. 

The aqueous acid dissolves iron, zinc, magnesium, and other 
metals with the formation of their chlorides and liberation of 
hydrogen. 

Aqua Regia. — A mixture of hydrochloric and nitric acids 
is known as aqua regi'a, because it dissolves the noble metals 
gold and platinum, the chlorides of these metals being formed. 
This is due to the liberation of chlorine and nitrosyl chloride, 
NOCl, during their mutual decomposition, thus : — 

HNO. + 3 HCl =: C\, + NOCl + 2 H,0. 



LESSON XVIII 



SOLUBILITY OF HCl 



179 



Experiments with Hydrochloric Acid Gas 




Fig. 84. 



ExPT. 98. — If a well-covered jar of the gas be inverted 
over water in the pneumatic trough, and the cover slipped off, 
the water rushes up and fills the jar 
owing to the great solubility of the 
gas. 

ExPT. 99. — The solubility of the gas, 
and also its acid nature, may be demon- 
strated by repeating Expt. 84, collect- 
ing the gas this time, however, by down- 
ward displacement, and using blue litmus 
solution, which is turned red on coming 
in contact with the gas (Fig. 84), show- 
ing that its reaction is acid. 

Expt. 100. — If the gas be led into 
a bottle half filled with water (Fig. 85), 
it will be seen that no bubble escapes, 
until the liquid becomes nearly saturated, 
the whole being absorbed by the water, 
giving rise to the aqueous acid. It is this aqueous solution 
which is commonly known as hydrochloric acid or muriatic 
acid, or sometimes as spirits of salt. This 
compound is a very pow^erful acid, and is 
much used in the arts. 

Determination of Composition of 
Hydrochloric Acid 

Expt. ioi. Electrolysis of the Acid. 

— The strong aqueous acid is decomposed 
into its elements, hydrogen and chlorine, by 
the electric current, but for this purpose 
platinum electrodes cannot be used as they 
were in the case of water, as chlorine rapidly attacks this 
metal. Carbon electrodes are therefore used for the decom- 
position. Fig. 86 shows the apparatus used for the electrolysis 
of hydrochloric acid. This experiment must be made in a 
very dim daylight or better in gaslight, as in bright sunlight 




Fig. 85. 



i8o 



ELECTROLYSIS OF HCl 



or even good daylight the mixed gases immediately combine 
with explosive violence. 

To ascertain the composition of the gases which are evolved 
by the electrolysis of the acid, it is necessary to allow the 
decomposition to go on for some time before collecting them, 

as the chlorine is 
soluble in the liquid. 
After a steady evolu- 
tion of the gas has 
continued for some 
time, the liquid be- 
comes saturated with 
the gas, and no fur- 
ther absorption takes 
place. The mixed 
gases may then be 
collected for the pur- 
pose of ascertaining 
their composition. 

For this object 
the mixed gases are 
allowed to pass 
through thin bulb 
tubes (Fig. 87), and 
after the gases have 
been passing long enough to ensure that all the air is driven 
out from the bulbs, they are very carefully sealed before the 
blow-pipe, and the sealed bulbs preserved for use in a dark box. 
ExPT. 102. Action of the mixed Gases on Potas- 
sium Iodide. — The capillary end of one of the bulb tubes is 
broken under mercury. It 
is seen that the bulb is 
full of the mixed gases 
and the mercury does not 
enter. Now, a solution of 
potassium iodide is placed on the top of the mercury, and 
the bulb tube raised so that a Httle of the solution enters. 
Iodine is now liberated by the chlorine, which is soon entirely 
absorbed, when it is seen that half the original vohcme of gas 
re7fiains, the iodine having dissolved in the liquid (Fig. 88). On 




Fig. 86. 




Fig. 87. 



LESSON XVIII 



COMPOSITION OF HCl 



iSi 



depressing the bulb in water and breaking the other end of the 
tube it will be found that the remaining gas takes fire and 
burns with the characteristic flame of hydrogen. The mixed 
gas then consists 

(i) of I volume of chlorine which is absorbed by the potas- 
sium iodide, the solid iodine dissolving in the liquid. 

KI + C1 = KC1 + I. 

(2) of I volume of hydrogen which remains and which may 
be burnt, giving the flame of burning hydrogen. 

ExPT. 103. Slow Combination of the mixed Gases 
in diffused Daylight. — If another bulb be taken from the 
dark box in which they must be kept, and 
exposed first to a dim, and afterwards gra- 
dually to a stronger daylight, the yellow 
colour of the chlorine will be noticed slowly 
to disappear, after wdiich the bulb may be 
exposed to strong daylight so as to com- 
plete the combination of the chlorine and 
hydrogen. If now the capillary tube of the 
bulb be broken under mercury, no change 
in the volume of the gas can be noticed, 
for no gas escapes and no mercury enters 
the bulb. If water be next poured on the 
surface of the mercury, and the end of the 
tube raised into it, the hydrochloric gas will 
be completely absorbed, and the water will entirely fill the bulb, 
showing that equal vohunes of hydrogen and chlorine combine 
together without change of volinne to form hydrochloric acid gas. 

H2 + CI, =2 HCl. 

2 vols . + 2 vols . = 4 vols . 




ExPT. 104. Explosive Combination in bright Daylight. 

— If a bulb filled with the mixed gases be taken from its dark 
box and exposed to bright daylight, immediate combination 
takes place throughout the mass, and so much heat is evolved, 
and the gas is thereby so greatly and suddenly expanded, that 
the bulb explodes and is shattered to fine dust. 

ExPT. 105. Explosion in the Light from burning 



l82 



CHLORINE AND HYDROGEN BULBS 



Magnesium. — A convenient manner of showing the explosive 
combination of the two gases is shown in Fig. 89. The bulb 
is exposed to the light from burning magnesium, the face and 
hand being protected by a sheet of plate glass. This light is 
rich in chemically active rays, and causes the combination just 
as daylight does. 

ExPT. 106. Action of Hydrochloric Acid on Sodium 
Amalgam. — Another experiment showing that hydrochloric 
acid gas contains half its volume of hydrogen may be made as 
follows. Into the open limb of the U tube (Fig. 90), mercury 




Fig. 89. 

is poured until both limbs are full, the air being driven out of 
the left-hand limb through the open stopcock, the stopcock is now 
connected with a tube delivering pure and dry hydrochloric 
acid gas. The pinch-cock at the bottom of the (j tube is now 
opened, so as to allow mercury to flow out, when the pure gas 
enters and fills the left-hand tube. After adjusting the level 
of mercury to the same height in both tubes, so that the gas 
fills the tube to the second mark (Fig. 90), sodium amalgam* 

* Made by pressing, by means of a pestle, several small pieces of clean- 
cut sodium, one by one, under the surface of a few ounces of mercury con- 
tained in a porcelain mortar. 



LESSON xviii ACTION OF SODIUM-AMALGAM 



183 



is poured into the open limb of the U tube so as to fill it 
completely. The open end of the tube is now closed by the 
thumb, and it is inclined and shaken so that the gas comes 
into intimate contact with the sodium amalgam. The gas is 
now again transferred to the closed limb, and the thumb 
removed from the open end, when a great diminution in the 
volume of the gas will be noticed. On adjusting the level of 
mercury to the same height in both tubes, it will be seen that 
the gas occupies just half the original volume (Fig. 91), whilst 
if mercury be poured into the open limb so as to fill it, the 
gas may be ignited on opening the stopcock^ and seen to burn 




Fig. 90. 



with the characteristic flame of hydrogen. The sodium of the 
amalgam has combined with the chlorine and liberated hydrogen 
which occupies half the volume of the hydrochloric acid as is 
shown in the following equation, 



HCl + Na : 

2 vols. 



NaCl + H. 

I vol. 



Test for Hydrochloric Acid. — Hydrochloric acid may 
be recognised by the fact that when heated with manganese 
dioxide chlorine is evolved, which may be recognised by its own 
characteristic tests. With silver nitrate solution, AgNOo, all 
soluble chlorides as well as the acid form a curdv white 



i84 THE CHLORIDES part ii 

precipitate of AgCl insoluble in nitric acid but soluble in 
ammonia. 

AgNOg + HCl = AgCl + HNO3. 



The Chlorides 

Occurrence. — The metallic salts of hydrochloric acid are 
called chlorides, they occur plentifully in nature ; common 
salt and rock salt, and also sylvine have already been men- 
tioned. Potassium chloride also occurs as a double chloride of 
that metal and magnesium in the Stassfurt deposits and is called 
Carnallite, KCl.MgCl2.6 HoO. Silver chloride, AgCl, occurs 
as horn silver, whilst ferric chloride is found in the craters of 
active volcanoes. The chlorides of the alkali metals, sodium 
and potassium, occur as essential constituents of the bodies of 
animals and plants. 

Preparation. — Chlorides may be formed either 

(i) by the direct combination of the metals with chlorine. 

Fe2 + 3 CI2 = Fe.Clg. 

(2) The replacement of the hydrogen in hydrochloric acid 
by the metals. 

Zn + 2 HCl = ZnCl, + H^. 

(3) By the combination of basic oxides and hydroxides 
with hydrochloric acid. 

NaOH + HCl = NaCl + H2O. 

(4) By the decomposition of carbonates by the acid. 

K2CO3 + 2 HCl = 2 KCl + H2O + CO2. 

(5) By the double decomposition of two salts which together 
produce an insoluble or difficultly-soluble chloride. 

Pb (N03)2 + 2 NaCl = FhCl, + 2 NaN03. 

This reaction is a type of an important class of reactions which 
are termed ^^ dotible decoinpositio7is ^ 

Before Pb|C1.2 Pb(N03)2 After PbC]2 PbCl j 

reaction (N03)2|h2 °% HQ reaction (N03)2H2 °^ 2 HNO3 



LESSON xviii BLEACHING POWDER 185 

In these reactions both sahs are mutually decomposed, and 
they may be considered as two couples, viz. the acid and basic 
parts of each, which change partners. The basic part of one 
combining with the acid part of the other, and vice versd. 



Separation of the Chlorides in Analysis 

When hydrochloric acid or a soluble chloride such as 
sodium or ammonium chloride, is added to a liquid contain- 
ing a large number of metallic salts in solution, a double de- 
composition takes place (as shown in 5, p. 184), causing the 
precipitation of the insoluble chlorides of silver, mercury (mer- 
Qwroiis chloride), and lead, 

AgCl HgCl PbClg. 

These three metals are, therefore, classed together as Group 
I. in qualitative analysis. The precipitated chlorides may be 
filtered from the solution containing the remaining metals, 
whose chlorides are soluble. Groups II. and III. will be 
mentioned under " The Sulphides " (p. 197). 



Bleaching Po"wder 

When chlorine gas is passed over slaked lime, Ca(0H)2, it 
is absorbed, and bleaching powder is produced, which is con- 
sidered to be a mixture of two compounds, viz. calcium chloride, 
CaCU, and calcium hypochlorite, Ca(0Cl)2, corresponding to 
hypochlorous acid, HCIO. The formation of bleaching powder 
may therefore be expressed as follow^s : — 

2 Ca(OH), + 2 CI2 = CaCl^ + Ca(0Cl)2 + 2 H2O. 

It is to the presence of the hypochlorite that bleaching 
powder owes its bleaching properties. For this compound is 
decomposed by hydrochloric acid, wdth liberation of chlorine, 

Ca(0Cl)2 + 4 HCl = CaCl2 + 2 H2O + 2 CI2. 

With sulphuric acid the reaction is rather different, although 
the result is the same. The sulphuric acid liberates the hydro- 



i86 SUMMARY AND EXERCISES PART ii 

chloric and hypochlorous acids from their salts, and the free 
acids mutually decompose with liberation of -chlorine and water, 
thus : — 

(i) CaCl2 + H2SO4 = CaS04 + 2 HCl 

(2) Ca(OCl)2+HoS04 =CaS04+2HC10 

(3) 2HQ + 2HC10 = 2H20 +2CI2. 

What we have learnt 

In the eighteenth Lesson we have learnt the various methods of pre- 
paring chlorine and hydrochloric acid, and their most important properties. 
We have seen that chlorine is a heavy, yellow, poisonous gas which dis- 
solves in half its bulk of water to form chlorine water. It combines directly 
with metals, bleaches organic colouring matters by an oxidising reaction 
in which water is decomposed and nascent oxygen liberated. Chlorine 
has a great affinity for hydrogen, and is able to abstract that element from 
many compounds containing it, such as H2O, H2S, C2H4, turpentine, etc. 
Equal volumes of chlorine and hydrogen form a mixture which only requires 
to be brought into strong daylight in order to explode. 

The composition of hydrochloric acid has been exhibited by various 
experiments, the mixture of gases evolved on the electrolysis of the aqueous 
acid being shown to consist of equal volumes of chlorine and hydrogen. 
The separation of chlorides as the first group in the qualitative analysis, 
of solutions of metallic salts has been explained, as well as the various 
methods of preparing chlorides. By passing chlorine over slaked lime we 
obtain bleaching powder, which may be considered as a mixture of cal- 
cium chloride and calcium hypochlorite. 

Exercises on Lesson XVIII 

1. By what reactions may chlorine be prepared ? Sketch the apparatus 
you would use, and give an account of the principal properties of chlorine 

2. What volume of chlorine can I get at NTP from 50 grams of man- 
ganese dioxide which is decomposed by an excess of hydrochloric acid ? 

3. Explain the bleaching action of chlorine. 

4. How is hydrochloric acid manufactured on the large scale ? Give 
equations for the reactions. 

5. How would you ascertain the composition of hydrochloric acid ? 

6. By what reactions may chlorides be prepared ? 

7. Explain the term " double decomposition." 

8. How are the metallic chlorides used in qualitative analysis ? 

9. W^hat is aqua regia ? 

10. Give characteristic tests for chlorine, hydrochloric acid, and 
chlorides. 

11. Explain the action of chlorine on (i) water, (2) sulphuretted hydro- 
gen, and (3) potassium iodide. 



LESSON XIX 

SULPHUR, SULPHURETTED HYDROGEN, AND THE 
SULPHIDES 

Sulphur 

Symbol S. Atomic Weight 31.8. Vapour Density 15.9 

Occurrence. — The ancients ^Yere acquainted with sulphur. 
It is found in the free or " native "' state near active as well as 
extinct volcanoes. The common name of brimstone is derived 
from Brennestone, or burning-stone, indicating its combus- 
tibility. Sulphur is not only found in the free state, but it is 
also met with in metallic ores, which contain sulphur combined 
with metals to form sulphides. Sulphur also occurs, combined 
with oxygen as well as with a metal to form sulphates, some of 
which occur in large quantities. Some of the more important 
sulphides amongst the metallic ores may be here mentioned : 
e.g. Galena, lead sulphide, PbS ; Cinnabar, mercury sulphide, 
HgS ; Iron pyrites, iron di-sulphide, FeS^, this, however, has 
no value as an iron ore, but as a sulphur ore ; Zinc Blende, zinc 
sulphide, ZnS. 

Among the commonly occurring sulphates are Gypsum, cal- 
cium sulphate, CaS04.2 H^O, this, w^hen heated, so as to drive 
off its 2 molecules of water of crystallisation, forms Plaster 
of Paris, CaSO^ ; Heavy spar, barium sulphate, BaSO^ ; 
Green Vitriol, ferrous sulphate, FeSO^.7 H.^O : Glauber salt, 
Na,SO^.ioH,,0. 

Sulphur is also found in nature in combination with hydro- 
gen, as sulphuretted hydrogen, H.,S. This gas occurs dissolved 

187 



NATURAL DEPOSITION OF SULPHUR 



in the water of many mineral springs, such as that at Harrogate, 
imparting to it the well-known smell of rotten eggs. 

Sulphur also occurs in a state of combination in many organic 
compounds, especially in animal substances, such as white of egg 
(albumen). When the albumen undergoes putrifaction, sul- 
phuretted hydrogen is formed, and this is easily recognised by 
its peculiar and unpleasant smell. 

Natural Deposition of Free Sulphur. — Volcanic gases 
contain sulphur in the form of sulphur dioxide, 80^, as well as 
of sulphuretted hydrogen, HgS ; and when these two gases 
come together, they mutually decompose with deposition of 
sulphur, and it is probable that this is the explanation of the 
occurrence of free sulphur in volcanic districts. Their mutual 
decomposition is shown in the following equation : — 

SO2 + 2H2S =38 + 2H2O. 

ExPT. 107. — This reaction can be readily shown by passing 
the two colourless gases into a large glass globe, when the 
deposit of solid yellow sulphur on the inside of the glass globe 
is clearly seen. 

Purification of Crude Sulphur. — In order to obtain pure 
sulphur from the impure material found in quantity near Etna, 
in Sicily, the crude sulphur, w^hich is found mixed with earthy 
and mineral matter, is placed in a kind of kiln built on sloping 
ground, so that, when fire is applied to the heap, a portion 
of the sulphur takes fire, and is burnt to sulphur dioxide, 
8O2, whilst the remainder, and by far the larger quantity, 
is melted and runs out at an opening at the lower part 
of the kiln, where it flows into moulds, and is then suffi- 
ciently pure for exportation. In order further to purify it, 
this crude commercial sulphur, when it reaches this country, is 
refined by subjecting it to distillation; as shown in Fig. 92. 
The sulphur is placed in an iron boiler (G) and heated by a 
fire, whilst the vapour (D) from the boiling sulphur passes 
into a cool brick chamber (A), where it collects. If the 
temperature of this chamber is kept below 115°, at which 
point sulphur melts, the vapour quickly solidifies in the 
form of a fine yellow crystalline powder, known as flowers, or 
flour, of sulphur, just as the vapour of water, if cooled below 
o°_, the melting point of ice, solidifies in the form of crystalline 



LESSON XIX ALLOTROPIC FORMS OF SULPHUR 



189 




snow. When the heat of the chamber rises above 115° the 
sulphur vapour liquefies, and the liquid is then drawn off and 
cast into moulds, when 
it is known as roll sul- 
phur or brimstone. 

Sulphur is now also 
largely manufactured 

from the " waste " of the 
alkali works, and this 
method is gradually 
superseding the produc- 
tion of native or volcanic 
sulphur. 

Allotropic Forms 
of Sulphur 

Sulphur is remark- 
able as existing in sev- 
eral distinct allotropic 
forms. The first of these 
is that in which it occurs free in volcanic districts ; it is there 
found crystallised in large transparent octahedrons belonging to 
the rhombic system shown in Fig. 93. The second form is also 
crystalline, and may be easily obtained as follows : — 

ExPT. 108. — Let us place some 
pieces of brimstone in a clay cru- 
cible, and heat the crucible in the 
fire, or over a gas flame, until the 
sulphur is completely melted. We 
then remove the crucible, and allow 
it to cool, until a thin crust of solid 
sulphur is formed on the surface of 
the molten mass. Two holes are 
then pierced through the crust, and 
the liquid sulphur in the interior 
quickly poured out. On then break- 
ing the crucible, or removing the 
mass of transparent needle-shaped 
These crys- 



Fig. 92. 




Fig- 93- 

whole of the top crust, a 

crystals is found to line the sides of the crucible. 



igb 



PLASTIC SULPHUR 



PART II 



tals are of a different shape from those occurring in nature, 
they belong to the monoclinic system of crystals. But if we 
allow this mass to remain in the crucible for twenty-four 
hours, it will then be seen that the crystals, which originally 
were transparent, have become opaque, and if they are then 
broken we notice that each needle-shaped crystal splits up 
into a number of small crystals, each of wdiich, under a micro- 
scope, is seen to have the form of the first modification. 
Hence, the first is the permanent, whilst the second is a less 
permanent form. These varieties differ not only in crystalline 
form but in certain other respects ; thus the specific gravity of 
the crystals of native sulphur is 2.05, whilst that of the crystals 
obtained by quickly cooling melted sulphur is 1.96. More- 
over, the first modification melts at 114°. 5, whereas the melting 
point of the second is 120°. 

ExPT. 109. — That sulphur exists in a third allotropic con- 
dition can be easily shown. For this purpose we melt some 
sulphur in a flask, and gradually heat it further, when the pale 
yellow mobile liquid changes to a dark-red viscid mass, 
until, when the temperature reaches 220°, the colour of the 
liquid is almost black. If I now pour out this thick liquid 
in a thin stream into cold water, the sulphur forms a soft 

sticky mass, something 
like caoutchouc, called 
plastic sulphur, because 
it can be drawn out into 
strings. Soon, however, 
it again assumes the 
yellow colour and brittle 
character of ordinary sul- 
phur. 

Fig. 94 shows the 
formation of plastic sul- 
phur from the distilled 
sulphur which has con- 
densed in a molten con- 
dition in the hot neck of 
the retort. 

Properties. — Sulphur is a yellow solid body, which when 
heated to 440^ in a retort begins to boil, giving off a red- 




Fig. 94. 



LESSON XIX TESTS FOR SULPHUR 191 

coloured vapour, which condenses in the neck of the retort, 
and sulphur may thus be purified on a small scale. 

When sulphur is heated in the air, it takes fire and burns 
with a blue lambent flame ; when burnt in oxygen (see Expt. 
36) the combustion proceeds more quickly, and the light and 
heat evolved in a given time are greater. In both cases the 
compound formed is sulpher dioxide, SO2, a colourless gas, 
sometimes called sulphurous acid gas, possessing the peculiar 
and well-known suffocating smell of burning sulphur. 

Sulphur does not dissolve in water ; it does so slightly in 
alcohol, and very easily in carbon disulphide, CS2, from which 
solution it crystallises in the same form as native sulphur. 

Sulphur combines "with Metals to form Sulphides 

Expt. no. — Instead of taking sulphur and iron filings as we 
did in Expt. 13, let us heat a mixture of flour of sulphur and 
finely divided metallic copper in a test-tube ; the mass will soon 
be seen to become red-hot, and, after cooHng, the tube is found 
to contain a black mass of copper sulphide, CuS. 

Expt. hi. — Another way of showing this same combination 
is to boil some sulphur in a flask, and then to lower a coil of 
copper wire into the dark vapour. The copper soon becomes 
red-hot, and the molten sulphide produced, drops to the bottom 
of the flask. 

Tests for Sulphur 

(i) Free sulphur may be recognised by the peculiar smell 
observed on burning it. 

(2) A more delicate test for sulphur, whether in the free 
state or in a state of combination, is by heating the 
substance with concentrated nitric acid, which oxidises 
the sulphur to sulphuric acid. This may be recog- 
nised in the solution by adding barium chloride, BaCU, 
which produces a heavy white precipitate of barium 
sulphate, BaSO^. When substances are present (such, 
for example, as barium carbonate) which render the 
sulphuric acid insoluble, a different method must 
be used, see ^'The Sulphates,'' p. 213. Further tests 



192 SULPHURETTED HYDROGEN part II 

for sulphur, when present as sulphides, sulphites, or 
sulphates, will be mentioned under their respective 
headings. 

Sulphur resembles Oxygen in its power of 
Combination 

We know (see p. 79) that carbon burns in the air, or in 
oxygen, to form carbon dioxide, CO^ ; now, although carbon 
does not burn in the same w^ay in sulphur vapour, yet, if we 
pass this vapour over red-hot charcoal, the charcoal gradually 
disappears, and we find that a colourless volatile liquid is formed 
by the union of the two elements. This liquid is carbon disul- 
phide, CSg, corresponding in composition to CO2. So, too, 
H^S, sulphuretted hydrogen, corresponds to HgO, water, whilst 
the sulphides of the metals correspond to the oxides. This 
analogy is well seen in the following list of some compounds 
of oxygen in the upper line, and of the corresponding sulphur 
compounds in the lowxr : — 

H,0 
HoS 

Compounds of Sulphur and Hydrogen 

Just as oxygen forms two compounds wdth hydrogen, viz., 
H^O and H^^O^, so sulphur forms two compounds, having a 
similar composition, viz., H^S, sulphuretted hydrogen or hydro- 
gen sulphide, and H2S2, hydrogen disulphide. The first of 
these is a gas at ordinary temperatures, the second is a liquid. 



Sulphuretted Hydrogen, 
Hydrogen Sulphide, or Hydrosulphuric Acid 

Formula H2S. Molecular Weight 33.8. Density 16.9 

Occurence. — It has already been stated that this com- 
pound is found dissolved in certain mineral waters, and also 
that it exists in volcanic emanations, and is produced by the 
putrefaction of animal matters, such as albumen, which con- 



CO., 


P.,0, 


K,0 


KHO 


CaO 


PbO 


HgO 


FeO 


cs^ 


P2S5 


KoS 


KHS 


CaS 


PbS 


HgS 


FeS 



LESSON XIX 



PREPARATION OF H.,S 



193 



tain sulphur in combination. It also occurs in the manufacture 
of coal gas, and has to be removed before the gas is sent out 
for consumption. 

Preparation. — Sulphuretted hydrogen is best prepared by 
acting with an acid on a metallic sulphide, and for this purpose 
sulphide of iron, FeS (obtained by heating together sulphur and 
iron filings, see Expt. 13), is generally employed. This sul- 
phide is acted upon by dilute sulphuric or hydrochloric acid. 
when the gas is rapidly evolved, sulphate or chloride of iron 
being formed, according to the kind of acid used. The reactions 
are as follows : — 



FeS + H,SO, = FeSO, 4- H,S, 
FeS + 2 HCl = FeCl. + H."s. 



Expt. 112. — The apparatus shown in Fig. 95 serves to pre- 
pare and purify the gas : the iron sulphide is placed in the 
flask provided with a thistle funnel, and dilute acid added : the 
gas evolved bubbles through the water contained in the second 
flask, and is thus purified from any acid which might pass or 
spurt over. The gas delivery tube, instead of being bent at 
right angles as in the Fig. 95, may be so arranged as to pass 
into a pneumatic trough, if it is desired, to collect jars full 
of the gas. The 
water in the trough 
and jars should, 
however, be warm, 
as water dissolves 
three times its 
volume of the gas 
at ordinary tem- 
peratures, but the 
solubility de- 
creases as the tem- 
perature of the 
water rises. It is 
advisable for laboratory use, as we shall see directly, to possess 
a means of obtaining either a small or large quantity of this gas 
at will. This is accomplished by using a Kipps's apparatus shown 
in Fig. 96. It consists of three glass globes, the two lower ones 




194 



KIPPS'S APPARATUS 



PART II 



(a) and (d) being connected by a narrow neck^ whilst the tubulus 
of the third globe (c) passes air-tight through the neck of (^) . 
Iron sulphide in lumps is placed in globe {d), and dilute sul- 
phuric acid poured in through the funnel tube until the globe (a) 
is filled, and some of the acid rises on to the sulphide of iron. 
Sulphuretted hydrogen gas is then evolved, passing out by the 
stopcock (e) into the wash bottle. When it is wished to stop 
the current of gas the cock (e) is closed, and in consequence of 
the gas accumulating in the globe (d)^ the pressure inside this 

globe increases, and thus 
forces out the acid up the 
tubulus into the upper globe 
(c)y then the evolution of gas 
ceases and the whole appa- 
ratus may be allowed to 
stand until the gas is again 
required. 

"Why the Gas thus ob- 
tained is not Pure. — 
If we wish to obtain perfectly 
pure sulphuretted hydrogen, 
we must use natural sulphide 
of antimony, Sb2S3, instead 
of artificial sulphide of iron, 
because this latter always 
contains some particles of 
metallic iron, and these, 
when acted upon by an acid, 
dissolve with evolution of 
hydrogen gas, and this 
hydrogen cannot be readily 
separated from the sulphuretted hydrogen. On the other 
hand, antimony sulphide, being a pure compound, dissolves 
without evolution of any hydrogen. Thus : — 

Sb.So + 6 HCl = 2 SbClg + 3 H.S. 

(2) Sulphuretted hydrogen is also formed in small quantities 
when hydrogen is passed through boiling sulphur. 

Properties. — Sulphuretted hydrogen is a colourless in- 
visible gas, which possesses a sweetish taste, and a very power- 




LESSON XIX PROPERTIES OF H^S 195 

ful and disagreeable smell, like that of rotten eggs. It is a 
poisonous gas, and, when inhaled in the pure state, it produces 
insensibility ; and, if the inhalation be continued, it causes 
death. Hence experiments with this gas should be made in 
a draught chamber or in the open air. 

ExPT. 113. — Collect two cylinders full of sulphuretted hydro- 
gen in the pneumatic trough over hot water. Apply a lighted 
taper to the mouth of one cylinder, and observe that the gas 
burns with a blue lambent flame, the hydrogen burning to 
form water, and the sulphur partly burning to sulphur dioxide 
(SO2), which can easily be recognised by its pungent smell, and 
partly being deposited as a yellow film on the sides of the jar. 

H.S + O = H2O + S, 
H^S + 30= HgO + SO^. 

The flame of a lighted taper plunged into the second cylinder 
is seen to be extinguished, and can be re-ignited at the mouth 
where the sulphuretted hydrogen is burning. 

Liquefaction of Sulphuretted Hydrogen. — Like all 
gases, this compound can be liquefied by cold or pressure, or 
both combined. If cooled to -62° it condenses to a colour- 
less liquid, which, if further cooled to -85^, freezes to an ice- 
like soHd. So, too, if we expose the gas to a pressure of about 
17 atmospheres, it likewise liquefies. 

Tests for Sulphuretted Hydrogen and the Sulphides 

(i) Sulphuretted hydrogen may be readily detected by its 
characteristic smell. 

(2) A piece of filter paper, moistened with a solution of 

lead acetate, is stained black when brought in contact 
with this gas, owing to the formation of black lead 
sulphide, PbS. 

(3) Sulphuretted hydrogen, or any soluble sulphide, gives a 

very characteristic purple colour to an alkaline solu- 
tion of sodium nitropmsside. 

(4) Sulphides are decomposed by hydrochloric acid, with 

liberation of sulphuretted hydrogen, which may be 
tested for by tests i and 2 above. 



196 COMPOSITION OF H2S part u 

From test No. 2 it is easily understood why white-lead 
paint IS soon turned black near gas-works,, sewers, or other 
places where sulphuretted hydrogen is allowed to escape into 
the air. 

The blackening of silver articles, and especially silver egg- 
spoons, is also due to the formation of black silver sulphide, 
Ag.S. 

Use of Sulphuretted Hydrogen in Analysis. — Sul- 
phuretted hydrogen is largely used in analysis for the sepa- 
ration of the sulphides of the metals of Group II. which 
are insoluble in acids, and for the further separation of the 
metals of Group III. which are insoluble in alkaline and 
neutral solutions. For further details see "The Sulphides,^' 
p. 197. 

Determination of Composition. — When strongly heated, 
this gas decomposes into sulphur and hydrogen, the sulphur 
being deposited in the solid form. For the purpose of deter- 
mining the composition of sulphuretted hydrogen, w^e take a 
tube, having sealed through the glass, a spiral of thin platinum 
wire, which can be heated white-hot by a current of electricity. 
Fill the tube with sulphuretted hydrogen up to a given mark, 
and place the tube over hot w^ater. Then pass a current of 
electricity through the coil of wire, and observe that sulphur 
is deposited. After a time allow the w^iole to cool, and notice 
that the bulk of gas is unchanged. Now remove the tube and 
light the gas, which burns wdth the blue flame of hydrogen, 
without formation of sulphur dioxide, showing that all the sul- 
phuretted hydrogen has been decomposed. What do we con- 
clude from this? (i) That sulphuretted hydrogen contains its 
own volume of hydrogen ; therefore, two volumes of the com- 
pound gas contain t\vo volumes of hydrogen weighing 2. But 
by weighing the sulphuretted hydrogen gas, chemists have 
found that it is 16.9 times as heavy as hydrogen, or two 
volumes weigh 33.8, when the same volume of hydrogen 
weighs 2. If we now deduct this w^eight of hydrogen from 
33.8, we have left 31.8 for the weight of the sulphur con- 
tained in two volumes of sulphuretted hydrogen, or the gas 
is composed of one atom of sulphur weighing 31.8, and two 
atoms of hydrogen weighing 2, and its formula is, therefore, 
H.,S. 



LESSON XIX THE SULPHIDES 197 



The Sulphides 

Occurrence. — Many sulphides occur in nature as metallic 
ores, and have already been mentioned under sulphur, e.g. 
Galena, Cinnabar, Zinc Blende, Iron Pyrites. 

Separation of the Sulphides in Analysis 

The filtrate which was obtained from the chlorides of Group 
I. of the metals (see p. 185), contains salts of the metals of all 
the succeeding groups. If sulphuretted hydrogen gas be passed 
through this solution, to which must be added free hydrochloric 
acid, a precipitate is formed, consisting of the sulphides of the 
metals of Group II., viz. mercury,* lead,t bismuth, copper, 
cadmium, antimony, arsenic, and tin; their formulae are — 

HgS, PbS, BigSg, CuS, CdS, Sb^S., As.S^, SnS. 

All these sulphides are insoluble in dilute hydrochloric acid. 
If they are filtered off, and the clear filtrate is made alkaline 
with ammonia so as to form ammonium sulphide with the free 
sulphuretted hydrogen dissolved in the liquid, a further pre- 
cipitation of the sulphides of the metals of Group III. takes 
place, viz. sulphides of iron, aluminium, chromium, nickel, cobalt, 
manganese, and zinc — 

FeS, AlgSg,^ Cr2S3,t NiS, CoS, MnS, and ZnS. 

These sulphides are soluble in hydrochloric acid, and hence 
were not precipitated in Group II., when free acid was present. 
They are, however, insoluble in neutral or alkaline solutions, and 
may, therefore, be precipitated together by adding ammonium 
sulphide to their neutral or alkaline solutions, or by passing 

* Mercury forms two series of salts, viz. the mercurous and mercuric 
salts corresponding to mercurous oxide, Hg-iO, and mercuric oxide, HgO, 
e.g., mercurous chloride (calomel), Hg.2C]2, insoluble in water. Mercuric 
chloride (corrosive sublimate), HgCl.2, soluble in water. 

t Lead chloride is slightly soluble in cold water, hence a little lead 
chloride remains dissolved in the filtrate from Group I. and gives rise to a 
precipitate of insoluble lead sulphide in Group H. 

X The hydrated oxides of these two metals are in fact precipitated, as 
their sulphides are at once decomposed. 



198 



SEPARATION OF METALLIC SALTS 



PART II 



sulphuretted hydrogen through their solutions containing an 
excess of free ammonia. The filtrate from, the sulphides of 
Group III. contains metals whose sulphides are soluble both 
in acid and in alkaline solutions, such as barium, strontium, 
calcium, magnesium, sodium, and potassium. The first three 
of these are separated from the rest by adding ammonium car- 
bonate to the clear filtrate from Group III., when a precipitate 
of the carbonates of Group IV. is obtained, viz., BaCOg, 
SrCOy, CaCOg. Group V. contains the metals magnesium, 



potassium, sodium, and ammonium. 



Separation of Metals into Groups in Analysis 

Let us make, for example, a mixture of the nitrates (all 
nitrates are soluble in water) of silver, lead, mercury (mer- 
curic), bismuth, copper, iron, manganese, zinc, barium, and 
calcium. 

(i) Add hydrochloric acid to the solution and filter off the 
precipitated chlorides. 



Residue (Chlorides of Group I.) 
AgCl, PbCla 



Residue (Sulphides of Group II.) 
HgS, PbS, BigSg, CuS 



Residue (Sulphides of Group III.) 
FeS, MnS, ZnS 



Residue (Carbonates of Group IV.) 
BaCOo, CaCO. 



Filtrate 
(2) Pass HgS for some minutes into 
the filtrate which has been acidi- 
fied with hydrochloric acid, filter 
off the precipitated sulphides 



Filtrate 

(3) Add NH4OH until very alkaline 

and again pass HgS 

I 

Filtrate 

(4) Add ammonium carbonate 

(NHJXOs 

L 



1 

Filtrate 

(Group V.) 



In a similar manner the individual metals in all the groups 
may be tested for and separated from one another ; but as it is 
not our purpose in this book to study these separations more 
fully, a work on Qualitative Analysis must be consulted for 
further information on this subject. 



LESSON XIX HYDROGEN BISULPHIDE 199 



Hydrogen Bisulphide, HgSg 

This substance is an oily liquid obtained by pouring a solution 
of calcium disulphide into hydrochloric acid : — 

CaS^ + 2 HCl = CaCl. + H.S. 

This compound is not used in the arts, but it is interesting as 
having a composition corresponding to H^O^, and because, like 
this substance, it possesses bleaching properties. It easily de- 
composes into sulphuretted hydrogen and free sulphur. 



What we have learnt 

In our nineteenth Lesson we have learnt how sulphur is found in nature 
both in the free state and in metallic ores, as sulphides and sulphates. 
We have learnt the methods by which sulphur is purified, and also its 
principal properties. Sulphur exists in three allotropic modifications, and 
forms compounds with metals called sulphides which correspond in com- 
position to the oxides. The occurrence and methods of preparing sul- 
phuretted hydrogen as well as its properties, tests, and mode of determining 
its composition, have been considered. Its use as a i-e-agent in analysis, 
and also the method of separating the metals into groups have been 
described, and thus an idea has been obtained of the principles upon which 
qualitative analysis is founded. 



Exercises on Lesson XIX 

1. By what reaction is native sulphur formed in volcanic districts ? 
Describe the various allotropic forms of sulphur. 

2. Name some naturally occurring sulphides and sulphates and give 
their chemical names and formulae. 

3. Explain the following reaction — FeS + H2SO4 - FeS04 + H2S. If 
I want ten litres of sulphuretted hydrogen (NTP), how many grams of 
pure ferrous sulphide must I employ ? 

4. Explain precisely how the composition of sulphuretted hydrogen has 
been determined. 

5. Give the tests for sulphur, sulphuretted hydrogen, and a sulphide. 

6. Explain in detail how the metals contained in a mixture of metallic 
salts may be separated into groups. 



LESSON XX 

OXIDES AND OXY-ACIDS OF SULPHUR 

SULPHUR DIOXIDE, SULPHUROUS ACID, AND THE 
SULPHITES. SULPHUR TRIOXIDE, SULPHURIC ACID, 
AND THE SULPHATES 

Sulphur Dioxide 
Sulphurous Acid Gas or Sulphurous Anhydride 

Formula SO2. Molecular Weight 63.6. Density 31.8 

Sulphur dioxide is a colourless suffocating gas, and is 
the compound to which burning sulphur owes its characteristic 
odour. The ancients knew that when sulphur burns, pungent 
smelling vapours are given off, and these fumes were then, as 
now, employed as a means of fumigation. Priestley, however, 
first prepared the pure gas in 1775. Its occurrence in volcanic 
gases has already been mentioned. 

Preparation. — (i) For laboratory purposes this gas is pre- 
pared by the action of hot concentrated sulphuric acid upon 
copper turnings. 

Cu + 2 H2SO4 = CuSO^ + 2 H.O + SO2. 

ExPT. 114. — The apparatus used is the same as for the prep- 
aration of chlorine or hydrochloric acid (Fig. 81). Copper 
turnings are placed in the flask so as to well cover the 
bottom. Cold concentrated sulphuric acid is then poured 
down the thistle funnel until the copper is covered. The flask 

2Q0 



LESSON XX SULPHUR DIOXIDE 201 

is then heated carefully on a sand-bath until the evolution of 
gas begins, when the heat must be moderated so as to keep 
the current of gas under control. Collect several jars of the 
gas in the draught chamber by downward displacement as it is 
very soluble in water, and is more than twice as heavy as air. 

(2) On the large scale, sulphur dioxide is prepared by 
burning sulphur or iron pyrites, FeS^, and immense volumes 
of the gas thus obtained are daily used in the manufacture of 
sulphuric acid (see p. 207) . 

(3) Pure sulphur dioxide may also be obtained by heating 
sulphur with concentrated sulphuric acid, the latter giving up a 
portion of its oxygen which oxidises the sulphur. 

S + 2 H2SO4 = 3 SO, + 2 H.O. 

(4) Charcoal is similarly oxidised to CO9, whilst the sulphuric 
acid is reduced as before to sulphur dioxide, 

C + 2 H2SO4 = 2 SO2 + CO, + 2 H2O. 

(5) The gas may also be obtained by the action of dilute 
sulphuric acid on sulphites (see p. 205). 

Na^SOg + H^SO^ = Na^SO^ + H,0 + SO,. 



Experiments -with Sulphur Dioxide 

ExPT. 115. — Repeat Experiments 98 to 100 with this gas 
instead of hydrochloric acid, proceeding exactly as before, to 
show its very great solubility in water and its acid properties. 

Properties. — Sulphur dioxide is a colourless gas possess- 
ing a very pungent suffocating odour which is well known as 
that given off by burning sulphur. It is 2.2 times as heavy as 
air and 31.8 times as heavy as hydrogen. It is very soluble 
in water, one volume at 0° absorbing nearly 80 volumes of 
the gas, whilst even at 20° nearly 40 volumes are dissolved, 
hence it cannot be collected over water. It can, however, 
be collected over mercury. It is a powerful antiseptic, and has 
been largely used in the preservation of meat, and is used 
for disinfecting purposes. It does not support the combustion 
of a taper. 



202 LIQUEFACTIOiN OF SULPHUR DIOXIDE part II 

Lead dioxide ignites when plunged into the gas. Lead 
sulphate being produced, 

PbO. + SOo = PbSO^. 

Sulphur dioxide (or sulphurous acid) like chlorine, possesses 
bleaching properties, but for another reason, i.e.^ because of 
its power of liberating nascent hydrogen from water, which 
reduces colouring matters to colourless compounds, whilst 
chlorine bleaches by an oxidising action. Sulphur dioxide 
is used for bleaching straw, silk, and woollen goods. 

SO. + 2 H.O = H.SG^ + H2. 

Liquefaction of Sulphur Dioxide. — This compound is 
the most easily condensible of all the gases, for it may be 




Fig. 97. 



10'' or a 



obtained in the liquid state at a temperature of —loP 
pressure at 0° of only 1.53 atmospheres. 

ExPT. 116. — To obtain the liquid the arrangement shown in 
Fig. 97 is employed. Sulphur dioxide is generated in the flask 
from a mixture of copper turnings and strong sulphuric acid, it 
is purified by passing through a small quantity of water con- 
tained in the wash-bottle, thence it passes into a glass 
worm-tube which is placed in a cylinder filled with a freezing 
mixture made by pounding up two parts of ice with one part 
of common salt (NaCi), by which a temperature of —18° is 



LESSON XX COMPOSITION OF SULPHUR DIOXIDE 



203 



reached. The gas here condenses to a colourless liquid which 
may be collected in a small flask also plunged into some of the 
freezing mixture. Liquid sulphur dioxide boils at —8°, and if 
cooled below —75° it freezes to a transparent solid. 

ExPT. 117. — Another simple experiment shows that sulphur 
dioxide gas can be liquefied at the ordinary temperature by 
pressure alone. For this purpose we take an ordinary strong 
glass tube closed and drawn out to a point at one end but 
open at the other. 
This is filled with the 
gas by displacement ; 
and then a plunger, 
made of an iron rod 
to the end of which 
is fitted a greased 
india-rubber stopper 
which fits the bore of 
the glass, is inserted 
and forced down the 
tube. When this 
plunger has been 
forced down so that 
the gas occupies 
about one-fifth of its 
original volume, it 
will be seen that drops 
of liquid are formed 
which run down into 
the drawn-out end of 
the tube. If the plun- 
ger be now quickly 
withdrawn the liquid 
is seen to boil, and so much heat is absorbed by the vaporisation 
of the liquid that a portion of it freezes to a white solid. 




Fig. 98. 



Determination of Composition of Sulphur Dioxide 

ExPT. 118. — How do we ascertain that the formula SO2 
represents the composition of this gas? For this purpose the 
apparatus shown in Fig. 98 is used. The bulb is closed by a 



204 SULPHUROUS ACID t>ART it 

hollow stopper through which two stout copper wires are 
cemented ; one of these ends in a small platinum spoon, the 
other in a thin piece of platinum wire which lies inside the 
spoon. A small piece of sulphur is placed over the thin wire 
in the spoon, and the bulb and tube having been filled with 
oxygen gas, and the stopper placed in position, and the level 
of mercury in the tubes observed, the sulphur is ignited by 
heating the wire by a current of electricity, care being taken 
to draw some mercury off by the tap so as to reduce the pres- 
sure of the gas and counteract the expansion due to the in- 
crease of temperature from the burning sulphur. As soon as 
the combustion is complete, the apparatus is allowed to cool, 
when it is found that the level of the mercury rises exactly to 
that which it occupied before the experiment, the mercury 
which had been allowed to flow out of the stopcock having 
been poured back through the open limb of the tube. What 
does this show? It shows that sulphur dioxide contains its 
own volume of oxygen. We know, however, from experiment 
that this gas is 31.8 times as heavy as hydrogen or two volumes 
weigh 63.6 when hydrogen weighs 2, but two volumes of oxygen 
weigh 31.8 or (15.9 x 2), therefore the weight of the sulphur in 
two volumes of this gas is 31.8 (one atom of sulphur) and that 
of the oxygen is also 31.8 (two atoms of oxygen), that is, sulphur 
dioxide is represented by the formula SO^, and is made up of 
equal weights of sulphur and oxygen. 

Sulphurous Acid, HgSOg 

ExPT. 119. — I^ instead of passing the gas SO^ into the coil 
tube shown in Fig. 97, we replace this coil by a straight tube 
and allow the gas to pass into a bottle of water, we shall 
notice that no bubble of gas reaches the surface of the water, 
it is all absorbed, one volume of water at the ordinary tempera- 
ture absorbing about 50 volumes of the gas. The solution thus 
obtained retains the characteristic smell of the gas ; it is an acid 
liquid, turning blue litmus paper red, and contains sulphurous 
acid, H2SO3, and only by prolonged boiling can all the SO2 be 
driven out of solution. If water at 3" be saturated with the gas, 
and the solution allowed to stand, a crystalline hydrate of sul- 
phurous acid is deposited. 



LESSON XX THE SULPHITES io^ 

The Sulphites 
Normal Sulphites, Acid Sulphites 

If an alkali, such as caustic soda, be added to sulphurous 
acid, the acid reaction disappears and a crystalline salt termed 
a sulphite is formed, thus : — 

H.SOo + 2 NaOH = Na.SO,, + 2 H,0. 

Sulphurous acid and caustic soda yield sodium sulphite and 
water. Sulphurous acid is the first example we have had of a 
dibasic acid, that is, it contains two atoms of hydrogen in the 
molecule, one or both of which may be replaced by a metal, so 
that two series of sulphites thus exist. In the first series only 
half the hydrogen is replaced, whilst in the second series all the 
hydrogen is replaced. 

H.SOg or ^>S03 NaHSOg or ^^>S03 Na.SOo or j^^>SO,. 

Sulphurous Acid Acid Sodium Sulphite or Normal Sodium 

(dibasic). • Sodium Bi-sulphite. Sulphite. 

pg^^^^S Calcium Bi-sulphite or Ca ^^ Calcium Bi-carbonate or 

^SO Bisulphite of Lime. ^COa Bicarbonate of Lime. 

The bisulphites are largely used as antiseptics. The formula 
of calcium bisulphite may be represented as above, for calcium 
is a dyad metal, replacing two atoms of hydrogen, and one 
atom of calcium may be supposed to decompose two molecules 
of the acid in order to replace two atoms of hydrogen. These 
acid salts are, as their name implies, both acids and salts, for 
they are compounds containing hydrogen replaceable by metals, 
i.e.^ acids, and they are also acids in which hydrogen has been 
replaced by a metal, i.e.^ salts. Carbonic acid, H^^COo (see p. 
228) is likewise dibasic ; its acid calcium salt may be represented 
as above. 

Sulphur Trioxide, or Sulphuric Anhydride 

Formula SO3. Molecular Weight 79.5. Vapour Density 39.75 

Preparation. — (i) Sulphur trioxide is a crystalline solid 
body, obtained by passing a mixture of perfectly dry sulphur 



2o6 SULPHUR TRIOXIDE PART 11 

dioxide and oxygen through a tube containing heated platinum 
sponge or platinised asbestos.* 

SO. + O = SO3. 

If the gases are passed over the platinum sponge when it 
is cold no combination takes place ; but directly it is heated, 
dense white fumes of the trioxide are produced, and if these 
are passed into a perfectly dry receiver, cooled by a freezing 
mixture, the trioxide condenses in the form of beautiful long 
white silky needles. To obtain these, every portion of the 
apparatus must be perfectly dry, as the compound combines 
with water so readily, that a trace of it causes the crystals to 
deHquesce with formation of liquid sulphuric acid. 

(2) Sulphur trioxide may also be prepared by the careful 
distillation of Nordhausen sulphuric acid or fuming oil of 
vitriol, which is concentrated sulphuric acid containing the 
trioxide in solution. 

(3) Sulphuric anhydride may also be prepared by abstracting 
the elements of water from anhydrous sulphuric acid by means 
of phosphorus pentoxide. 

Properties. — Sulphur trioxide is a crystalline solid, which 
melts at 14.8°, and boils at 46°. It absorbs moisture readily 
from the atmosphere, and evolves dense white fumes of sulphuric 
acid. It combines with water with a hissing sound, and with 
the evolution of a large amount of heat. It is an acid-forming 
oxide, and combines with the basic oxide. Baryta, BaO, with 
such force, that in the formation of the salt, BaSO^, the mass 
becomes red-hot. 



Determination of Composition 

On passing the vapour of the trioxide through a red-hot 
tube it is decomposed into two volumes of SO., and one 

* Platinised asbestos is made by igniting asbestos fibres which have 
been covered with the yellow crystalline double chloride of platinum and 
ammonium, P1C14.2(NH4)C1. This decomposes on ignition leaving be- 
hind metallic platinum in a finely divided state. Platinum sponge is made 
by igniting the double chloride alone. 



LESSON XX SULPHURIC ACID 207 

volume of oxygen, and since its vapour density is found to be 
39.75, its formula must be SO3. 

SO3 = SO^ + O. 

2 Vols. I Vol. 

Sulphuric Acid, Vitriolic Acid or Oil of Vitriol, 

Sulphuric acid is the most important and most useful of all 
the acids, and by its means nearly all the other acids are 
prepared. It is manufactured on an enormous scale, nearly a 
million tons being annually manufactured in Great Britain, 
thus giving rise to a most important modern industry. 

. Manufacture of Sulphuric Acid 

The manufacture of sulphuric acid on the large scale 
consists in allowing (i) sulphur dioxide; (2) steam; (3) air, 
and (4) a small amount of the fumes of nitric acid to pass 
together into a capacious chamber hned and floored entirely 
with sheet lead (Fig. 100). We may suppose for simplicity, 
that the sulphur dioxide and steam unite to form sulphurous 
acid, which is afterwards oxidised to sulphuric acid by the 
oxygen of the air. 

SO2 + H2O = H2SO3. H2SO3 + O = H0SO4. 

But, as a matter of fact, this oxidisation only takes place very 
slowly indeed if atmospheric air alone be used, and here it is 
that the action of the nitric acid vapour, which is of the 
greatest interest, must be taken into account. 

The vapour of nitric acid, made from Chili saltpetre, NaN03, 
and sulphuric acid (see p. 149) is passed continuously in 
small quantity into the flue, along which is passing the sulphur 
dioxide, obtained by burning either sulphur itself, or the 
sulphur contained in iron pyrites, FeS2. 

The nitric acid is at once reduced to nitrogen-peroxide, which 
in turn is further reduced to nitric oxide. Thus : — 

SO, + 2 HNO3 = H,S04 + 2 NO,, 
NO2 + SO2 + H2O = H2SO4 + NO. 



SULPHUkiC ACID MANUFACTURE^ 



PARt It 



The nitric oxide, as we know, will at once absorb oxygen from 
the air, and be converted to the red fumes of the higher 
oxides, N^Oo and NO^. 

NO + O = NO^. 

This peroxide in turn reacts on more sulphur dioxide and 
steam, to form sulphuric acid, nitric oxide again being formed. 
In this manner the nitric oxide which is obtained by the 
reduction of the vapour of nitric acid, acts as a carrier of 




Fig. 99. 

atmospheric oxygen^ alternately abstracting it from the air, and 
then giving it up again for the oxidation of the sulphur 
dioxide ; hence it is that a very small proportion of nitrous 
fumes will suffice to bring about an oxidation, by atmospheric 
oxygen, of very large volumes of sulphur dioxide, which 
atmospheric air alone would be quite powerless to effect. 

The reactions by which sulphuric acid is thus produced on 
a large scale may be illustrated by means of the apparatus 
shown in Fig. 99. Sulphur is heated in the bulb tube S, 



LESSON XK THE LEADEN CHAMBER 209 

and allowed to burn in a stream of air forced forward by 
a double aspirator. A mixture of nitre and sulphuric acid 
in a evolves nitric acid vapours which pass into the large 
flask along with the sulphur dioxide and air. Here the 
gases meet with steam from the boiling water in b^ and they 
condense on the sides of the large flask as drops of sulphuric 
acid. By alternately increasing and diminishing the supply of 
sulphur dioxide, the disappearance and reappearance of the red 
nitrous fumes can be readily shown. If the flask into which the 
gases are passed is kept dry by keeping back the steam, white 
crystals, known as leaden chamber crystals, are formed, having 
the formula HS0o(N02). 

2 SO, + H^O + N2O3 + 0, = 2 S0,<^^^ 

we may suppose it to be sulphuric acid in which one of the 
hydroxyl groups (OH) is replaced by NO,. Thus : — 

SO <'^^ SO <-^^ 

Sulphuric acid. White leaden-chamber crystals. 

When steam is admitted the crystals dissolve with formation 
of sulphuric acid and red fumes of N^O.^. 

2 HSO3.NO0 + 2 H.O = 2 HoSO, + N.Oo + H,0, 
or SO,<2o,^ +H.OH = SO,<g{J + HNO,. 

It is evident that in the manufacture of sulphuric acid on a 
large scale, it is of the utmost importance to adjust the due 
proportion of the various gases, viz. (i) sulphur dioxide, (2) 
air, (3) nitric fumes, and (4) steam, with great nicety. If too 
much steam be passed into the chambers, they become too 
hot, and the acid is, moreover, weakened. If the nitrous 
fumes are deficient, the oxidation of the sulphur dioxide does 
not take place properly. If too much air be admitted, the 
gases are unnecessarily diluted, and the reactions do not take 
place so readily. 

The gases which pass into the leaden chambers (Fig. 100) are 
drawn through them by a powerful draught produced by the 
uprush of heated air in a tall chimney. Besides the residual nitro- 



210 



PROPERTIES OF HoSO. 



PART n 



gen from the air, a certain amount of the nitrous fumes pass out 
with the nitrogen at the end of the chamber. ' These red fumes 




are absorbed by an arrangement called a Gay-Lussac tower 
down which cold and concentrated sulphuric acid is allowed to 
trickle slowly before the waste gases pass into the upcast shaft. 



LESSON XX EFFECTS OF HEAT 211 

The red fumes are used again in the initial stages of the process 
by allowing the nitrated sulphuric acid obtained from the Gay- 
Lussac tower, to trickle down another tower called the Glover 
tower, or denitrating tower, together with the weak chamber 
acid. The hot gases from the pyrites kilns are first allowed to 
pass into the Glover tower, and the sulphur dioxide not only 
robs the nitrated acid of its nitrous fumes, but the hot gases 
are themselves cooled before they enter the leaden chamber, 
and the weak acid is at the same time made stronger by being 
deprived of a portion of its water which passes into the chamber as 
steam, and assists in the subsequent formation of sulphuric acid. 

The acid, as made in the leaden chambers, called ^* chamber 
acid,^' is a dilute acid of sp. gr. i .6. To obtain a stronger acid, 
this is first evaporated in leaden pans until the sp. gr. reaches 
1.72, when the acid is known in commerce as BroAvn Oil of Vitriol 
or B.O.V. 'It cannot be further concentrated in leaden pans, 
because the stronger acid attacks the lead. The B.O.V. must, 
therefore, be further concentrated in platinum or glass vessels. 
The acid thus obtained is far from pure, not only does it 
contain small amounts of nitrous fumes and sulphur dioxide, 
but also arsenic derived from the iron pyrites, wdiich are burnt 
as a source of sulphur, and lead sulphate derived partly from 
the leaden chamber, but more particularly from the leaden 
concentrating pans. To obtain the pure acid the commercial 
B.O.V. must be distilled. The volatile impurities pass over in 
the first third of the distillate, after which the pure concen- 
trated acid may be collected in a fresh receiver. This acid is 
free from nitrous fumes, arsenic, and lead, and also from the 
organic matter which renders the impure acid brow^n. 

Properties. — The acid thus purified by distillation still 
contains about 2 per cent of water w^hich cannot be removed 
by this process. If, however, the distillate be cooled, the pure 
acid, H^SO^, separates out in the form of crystals, which melt 
at 10.5°. The strong acid is an oily liquid having a specific 
gravity of 1.854 at 0°. 

Effects of Heat. — When the pure acid is heated it begins 
to fume at 30°, owing to a partial decomposition into H2O and 
SOo. This decomposition increases until the boiling point 
338° is reached, when a liquid remains wiiich contains 98.4 to 
98.8 per cent of the pure acid which distils without change. 



PROPERTIES OF H2SO4 



When very strongly heated, as, for example, when the acid 
is allowed to fall, drop by drop, into a red-hot platinum flask, 
filled with pumice stone, it decomposes into sulphur dioxide, 
oxygen, and water. 

H2SO4 = H2O + SO2 + O. 

The first two may be absorbed by water and the oxygen gas 
collected. 

Aflinity for Water. — When sulphuric acid is mixed with 
water, great heat is evolved owing to their chemical combina- 
tion, and a contraction in volume takes place. So much heat 
is thus given off that it is dangerous to add water quickly to a 
large bulk of the acid, the two combining with explosive vio- 
lence. In diluting the acid, therefore, the strong acid should 
be gradually poured into the requisite volume of water contained 
in a thin glass flask, the mixture being shaken or stirred during 
the addition of the acid. Owing to its great affinity for water, 
strong sulphuric acid is used for drying gases. For this pur- 
pose the gas is best passed through a tube containing pumice 
stone which has been boiled in the strong acid. This method 
cannot, of course, be used for drying such gases as ammonia, 
which combine with the acid. Granular anhydrous calcium 
chloride must then be substituted. Strong sulphuric acid is 
also used for keeping a dry atmosphere in the closed space of 
desiccators, either at the ordinary pressure or in vacuo. Solid 
substances are enclosed in this dry space to be kept dry, or to 
be deprived of the last traces of moisture. 

Not only does strong sulphuric acid combine energetically 
with water itself, but it also abstracts the elements of w^ater from 
many organic compounds, thus : — 

Formic Acid CH2O2+ H2SO4 - H2SO4.H2O + CO. 
Oxalic Acid C2H2O4 + H2SO4 = H2SO4.H2O + CO + CO2. 
Alcohol C2H6O + H2SO4 = H2SO4.H2O+ C2H4 (Ethylene). 

We shall see (Expt. 120) that the acid abstracts the elements 
of water from such carbohydrates as sugar in a similar manner. 

Crystalline Hydrate of Sulphuric Acid. — Sulphuric 
acid forms a definite cr3^stalline compound with water, having 
the composition H2SO4.H2O. A mixture of the acid with 



LESSON XX THE SULPHATES 213 

water in equal molecular proportions solidifies to a mass of 

prismatic crystals at 7.5°. 

Action on Metals. — The concentrated acid does not act 

on many metals in the cold although it does so when heated. 

Copper, mercury, antimony, bismuth, tin, lead, and silver are 

attacked by the hot acid with evolution of sulphur dioxide, 

thus : — 

Ag2 + 2 H,SO^ = Ag,SO^ + 2 H.,0 + SO2, 
Cu^ + 2 H^SO^ = CuSO^ + 2 H2O + SO2. 

Gold and platinum are unacted upon even by the boiling 
acid, and this acid is therefore used for parting or separating 
silver and gold. Such metals as zinc, iron, manganese, and 
magnesium are dissolved by the dilute acid in the cold with 
evolution of hydrogen, and the formation of sulphates. 



The Sulphates 

Sulphuric acid, like sulphurous acid, is a dibasic acid, and 
therefore forms two series of salts, viz., the acid salts and 
normal salts, thus : — 

g>SO, ^='>S0, Na>SO,. 

Many sulphates are found in nature, e.g. Gypsum, 
CaS04.2 HoO, Heavy spar, BaSO^, Celestine, SrSO^, Glauber 
salts, Na2SO^.ioH20, Epsom salts, IVLgSO^^.y H^O. 

Barium and lead sulphates are insoluble in water. Calcium 
and strontium sulphates are slightly soluble, whilst most other 
sulphates are soluble. 

Tests for Sulphates. — (i) All soluble sulphates, on the 
addition of barium chloride, give a white precipitate of barium 
sulphate insoluble in dilute hydrochloric acid. 

Na^SO^ + BaCl^ = BaSO^ + 2 NaCl. 

(2) If a sulphate is mixed with sodium carbonate and heated 
on charcoal in the reducing blowpipe flame a sulphide is pro- 
duced. If the fused rnass is laicj on a clean silver coin and 



214 SUMMARY 



moistened with water, a black stain of silver sulphide is formed 
on the coin. 

(3) Insoluble sulphates, such as barium sulphate, may be 
decomposed by boiling with sodium carbonate, when soluble 
sodium sulphate is produced, which may be filtered and tested 
for with barium chloride after acidifying with hydrochloric 
acid. 

BaSO^ + Na,C03 = BaCOg + Na.SO^. 



What w^e have learnt 

In our twentieth Lesson we have learnt that sulphur dioxide may be 
prepared (i) by the action of hot concentrated sulphuric acid on copper; 

(2) by burning sulphur or iron pyrites in the air ; (3) by heating sulphur 
or charcoal with strong sulphuric acid ; (4) by the action of acids on 
sulphites. 

We have seen that sulphur dioxide is a colourless, pungent smelling 
gas, about twice as heavy as air, and very soluble in water, wdth which it 
combines to form sulphurous acid. We have seen how it may be liquefied, 
and how its composition may be determined. 

Sulphurous acid is a dibasic acid, and therefore forms two series of 
salts, viz. normal sulphites, M'2S03; acid sulphites, M'HSOg, where M' 
stands for any monad metal. 

Sulphur trioxide is prepared by passing sulphur dioxide and oxygen over 
heated platinised asbestos. It is a white solid compound, which rapidly 
absorbs water with formation of sulphuric acid. This latter compound is 
manufactured on the large scale by passing (i) sulphur dioxide, (2) steam, 

(3) air, and (4) a small amount of nitric acid vapour into a capacious 
leaden chamber, where the nitrous fumes act as a carrier of the atmo- 
spheric oxygen to the sulphur dioxide, which is thus oxidised to sulphuric 
acid in the presence of steam. Sulphuric acid is a heavy oily liquid 
possessing very great affinity for water, the component elements of which 
it will abstract from many compounds. Like sulphurous acid, sulphuric 
acid is dibasic, and forms two series of salts, M'2S04, M'HS04. Many 
of the normal salts are found in nature. All soluble sulphates give a 
white precipitate with a solution of barium chloride, insoluble in hydro- 
chloric acid. 

Exercises on Lesson XX 

1. How is sulphur dioxide prepared (i) in the laboratory, (2) on the 
manufacturing scale ? State its principal properties, and the method of 
determining its composition. 

2. What is the action of strong sulphuric acid on carbon, on sulphur, 
on oxalic acid, on alcohol, and on formic acid ? 



LESSON XX EXERCISES 215 



3. How may sulphur dioxide be liquefied ? 

4. What weight of sulphur is contained in 2000 litres of sulphur dioxide, 
measured at 17*^ R. and 785 mm. ? 

5. Give the formulas of the sodium, calcium, and ammonium salts of 
sulphurous acid. 

6. How is sulphur trioxide prepared ? What special precautions must 
be taken in making it ? 

7. Describe the manufacture of sulphuric acid. 

8. How would you obtain sulphuric acid free from lead, arsenic, and 
nitrous fumes from a sample of B.O.V. ? 

9. What are leaden chamber crystals, and what is the action of steam 
on them ? Give formulas and equations. 

, 10. Explain the action of the nitric acid vapour which is used in the 
manufacture of sulphuric acid. 

11. What weight of sulphuric acid may be obtained by burning two tons 
of sulphur, supposing it to be completely oxidised to sulphuric acid ? 

12. Give the tests by which you would recognise sulphur, sulphuretted 
hydrogen, a sulphite, and a sulphate respectively. 



LESSON XXI 

CARBON AND ITS ALLOTROPIC MODIFICATIONS — CAR- 
BON MONOXIDE — CARBON DIOXIDE — METHANE- 
ACETYLENE — ETHYLENE — COAL GAS AND FLAME 

Carbon 

Occurrence. — Carbon is an essential constituent of every 
animal and vegetable body, from the most minute and simple 
organism up to the largest and most complicated. 

The number of the known compounds of carbon with three 
elements, viz. hydrogen, oxygen, and nitrogen, is far larger 
than that of all the other elements put together ; moreover, 
many of them possess so complicated a composition and are 
so different in their properties from the compounds of other 
elements, that they are usually considered as forming a separate 
branch of the science, and this branch is called organic 
chemistry. Still the study of carbon and some of its simpler 
compounds cannot be omitted from even an introduction to 
Inorganic Chemistry., such as the present work, because carbon 
itself possesses most interesting properties, and its simpler 
combinations, which may be considered to be inorganic 
compounds, play so important a part in the economy of 
nature. 

Whilst carbon itself occurs only in small amount, its com- 
pounds are found to exist in large quantity. It exists combined 
with oxygen in the atmosphere as carbonic acid gas, or carbon 
dioxide, CO^, and although its relative proportion is small 
(only about 4 in 10,000 vols, of air), the absolute amount is 
enormous, reaching several billions of tons. Besides, combined 

216 



LESSON XXI CARBON 217 

with lime, CaO, as limestone, CaCO.., and as magnesian lime- 
stone or dolomite, (CaMg)C03, it forms whole mountain 
ranges ; whilst as chalk and marble it exists in very large 
quantities. Chalk and coral have been formed partly, during 
past geological ages, through the agency of minute organisms 
living in sea-water. In addition to these, other enormous masses 
of carbon compounds exist in the living forms, animal and 
vegetable, existing on the earth's surface. Combined with 
hydrogen, carbon exists in rock oil or petroleum, benzene and 
coal gas, whilst in further combination with oxygen it occurs in 
vegetable oils, fats, sugar, starch, and woody fibre. 

ExPT. 120. — A striking experiment shows that white 
crystalline sugar contains large quantities of carbon. Let us 
place an ounce of lump-sugar in a large cylinder, and pour 
upon it enough hot water to cover it, and then add about 
double the volume of strong sulphuric acid, stirring up the 
mixture with a glass rod. In a few seconds the colourless 
liquid will become brown, then black, and in a few minutes it 
will boil up, the jar becoming filled with a coal-black mass of 
carbon. Sugar is a carbohydrate, and the strong sulphuric 
acid has abstracted from it the elements of water, leaving 
only carbon. 



AUotropic modifications of Carbon 

Carbon, like sulphur, is remarkable as existing in three 
totally difi'erent soUd forms : (i) diamond, (2) graphite, (3) char- 
coal ; hence carbon is said to exist in three allotropic modifica- 
tions. No one has yet succeeded in converting graphite or 
charcoal into diamond. How, then, do we know that the 
colourless sparkhng gem, the hardest of all known substances, 
is the same chemical substance as soft black graphite, com- 
monly called blacklead, and used for making pencils and 
polishing our stoves, and that this latter substance is again 
the same as soft porous charcoal got by heating wood, or as 
soot or lampblack, which is nearly the same thing as char- 
coal ? No one could have believed that these three substances 
consist of carbon until experiment proved this to be the case. 
Let us see how this was done. 



2i8 ALLOTROPIC FORMS OF CARBON part ii 



The Diamond, the purest form of Carbon, 
Crystallised Carbon 

The diamond has for ages been valued as a precious stone 
on account of its brilliant lustre and its great hardness (ada- 
mant) ; but its composition remained unknown until about 
a hundred years ago. Experiments had indeed been previ- 
ously made, which showed that diamonds could be burned ; 
thus diamonds placed in the focus of a powerful burning-glass 
were found to disappear, but what became of them was not 
known. Then diamonds and rubies (another kind of precious 
stone) were heated together in a furnace, and the diamonds 
disappeared, whilst the rubies remained unaltered — this also 
could not be explained. Afterwards it was noticed that if the 
diamonds were heated in a perfectly closed vessel they did not 
disappear. No one was able to explain these facts until 
Lavoisier made the following experiment about the end of last 
century. He placed a diamond in a glass vessel containing 
pure air over mercury, and then heated the diamond by means 
of a burning-glass ; the diamond took fire and burnt com- 
pletely away, leaving only a mere trace of ash, and on examin- 
ing the air after the combustion he found that it turned lime- 
water milky, and, therefore, contained carbonic acid gas, 
which, as we have seen, is always formed when carbon itself, 
or a carbon compound, is burnt in the air or in oxygen. But 
it may properly be said, it is true, that this experiment shows 
that diamond contains carbon, but it does not prove that it 
consists of carbon and nothing else. This, however, was 
ascertained by Davy in 1814, because he observed that no 
trace of water is formed when a diamond is burnt in pure 
oxygen, so that this gem can contain no hydrogen. Then 
further experiments proved that if we take the same weights, 
say 12 parts, of diamond, of graphite, of charcoal, they all 
yield exactly the same weight (44 parts) of carbonic acid gas. 
So now we know that each of these three substances consists 
of nothing else but carbon. 

Carbon in the free state exists only in small quantities in 
nature. The diamond is a rare and valuable substance (see Fig. 
1 01), it crystallises in forms derived from the regular octohedron, 




LESSON XXI DIAMOND, GRAPHITE, CHARCOAL 219 

is found in a peculiar rock or in alluvial deposits in Brazil, 
South Africa, the Urals, and elsewhere. Its specific gravity 
is about 3.5. We do not know how diamonds have been 
formed, nor how to make them arti- 
ficially. 

Graphite. — Graphite, the only 
other form of naturally occurring free 
carbon, is a soft black substance, of 
specific gravity about 2.2, and is found 
in many localities, the most famous of 
which is Borrowdale in Cumberland, 
but the mines there are exhausted. 
It has also been found in Siberia, in 
many places in North America, espe- 
cially in California. It occurs in 

granite, gneiss, and other crystalline rocks : its mode of forma- 
tion is likewise unknown. It sometimes occurs in crystals, 
but not in the same form as diamond. 

Charcoal is the amorphous or non-crystalline form of 
carbon. It is obtained by the decomposition of organic bodies. 
Thus when w^ood is heated strongly in absence of air, charcoal 
is left behind as a black porous mass, the volatile portions of 
the wood escaping. There are several varieties of charcoal — 
they are (i) lampblack; (2) gas carbon; (3) coke; (4) ani- 
mal charcoal : and (5) wood charcoal. 

"Wood charcoal is made by arranging pieces of wood in a 
heap, covering the heap with soil, leaving an air hole in the 
centre, and then kindling the wood at the bottom, care must be 
taken only to allow a small amount of air to enter, otherwise 
the whole heap would be completely burnt and only white 
ashes left. Charcoal is very porous, and, though heavier than 
water (sp. gr. about 1.8), it floats on this liquid, because its 
pores are filled with air; but if we extract this air by means 
of an air pump, the pieces of charcoal will be seen to sink in 
water. This porous nature of charcoal is a valuable property, 
as it enables charcoal to absorb large quantities of gases and 
colouring matters. 

ExPT. 121. — Thus if a few bits of charcoal, which have 
been previously heated in a flame, be passed up into a tube 
filled with dry ammonia gas collected over mercury, it is soon 



CHARCOAL PART II 



seen by the rise of the mercury in the tube that the gas is 
absorbed. Charcoal can absorb 90 times its bulk of NH3. 
Other gases are also absorbed by charcoal, but in smaller 
quantity. Thus it absorbs 9 times its volume of oxygen, and 
this is the reason why charcoal acts as a powerful disinfectant. 
If a piece of animal flesh be buried in charcoal powder, no 
unpleasant smell is noticed, the putrefactive gases which are 
given off from decomposing animal matter are absorbed and 
oxidised to carbonic acid and water by the oxygen taken up 
by the porous charcoal. 

Animal charcoal or bone black is made by charring 
bones in iron cylinders ; it contains, together with charcoal, the 
inorganic constituents of the bone (phosphate of lime). This 
is largely used for decolouring raw sugar, the brown syrup is 
allowed to run over bone black contained in iron cylinders, it 
passes out at the bottom of the column perfectly colourless. 

ExPT. 122. — The decolourising power of animal charcoal 
is well shown by adding a little blue solution of indigo to some 
boiling water, and then shaking this up with some bone black. 
After a short time the liquid is passed through a filter when 
it is seen to be colourless ; some of the blue liquid, without 
addition of bone black poured on to another filter, is as darkly 
coloured after filtration as before. 

Lampblack is a pure form of soot, made by imperfectly 
burning turpentine, resin, tallow, oil, or pitch. It is used for 
making black paint and printers^ ink. Lampblack is not pure 
carbon, as it contains oily matter. To obtain pure carbon from 
it, the black must be not only strongly heated, but ignited in 
a current of chlorine gas, as this is the only way of getting 
rid of all the hydrogen ; as it combines wdth the chlorine 
to form hydrochloric acid, leaving pure carbon behind, for 
lampblack, carefully made, contains no inorganic or mineral 
matter. 

Coal is the result of the decay and gradual decomposition 
of vegetable matter once growing on the earth's surface. This 
has been going on for ages. Coal is not pure carbon; it 
contains not only the mineral matter which the vegetation 
contained, but much more which has been added since the 
vegetation grew. Besides, it contains more or less hydrogen, 
r.s may be seen from the blaze which ordinary coal gives in the 



LESSON XXI COAL 



fire. This part is made use of in the manufacture of coal gas 
for illuminating purposes ; no gas can be made by heating 
charcoal or coke in a retort. There are many varieties of 
coal. Cannel (or candle) coal is richest in hydrocarbons ; 
anthracite coal poorest. Coal also contains sulphur (about 
2 per cent on the average) generally as iron pyrites or '^ coal 
brasses," FeS^. When coal burns, this sulphur also burns, 
forming sulphur dioxide, and this oxidises in contact with the 
atmospheric oxygen, and combines with moisture, forming sul- 
phuric acid. This is very deleterious and harmful to vege- 
table life, and it likewise attacks and destroys stone buildings. 
Many hundreds of tons of this acid are thrown into the atmo- 
sphere in London, where millions of tons of coal are burnt every 
year. 

Carbon is called a reducing agent. — What do we mean 
by that? An experiment will show us. 

ExPT. 123. — Take a little black oxide of copper, and mix 
with it some powdered charcoal ; heat the mixture in a test 
tube fitted with a delivery tube, which dips under some clear 
lime-water. What is observed? Gas is given off, and this, 
passing through the lime-water, makes it milky. So carbonic 
acid gas is given off. What change has here occurred? If 
we examine the residue it will be seen to contain bright red 
particles of copper. The carbon has reduced the oxide — 
thus 2 CuO + C = 2 Cu 4- CO^. Both coal and coke are largely 
used in metallurgical operations for this reason. Iron ore is 
smelted with coal and cast iron is produced. Many other 
metals are manufactured from their ores by using carbon as a 
reducing agent. 



THE OXIDES OF CARBON, CO AND CO2 

Carbon Monoxide, or Carbonic Oxide Gas 

Formula CO. Molecular Weight 27.8. Density 13.9 

This gas is formed when carbon is burnt in a limited supply 
of oxygen. The blue lambent flame so often seen at the top 
of a red-hot coal fire is that of carbon monoxide formed by the 



CARBON MONOXIDE 



PART II 



action of the red-hot fuel on carbon dioxide produced by the 
combustion of the fuel at the lower part of the fire ; thus : — 

CO2 + C = 2CO. 

The carbon monoxide burns again to CO2 when it reaches the 
top of the fire, and there comes in-to contact with the oxygen of 
the air. This reaction serves as a means of preparing this 
compound. For this purpose a piece of iron gas-piping filled 
with charcoal and furnished with corks and tubes as shown in 
Fig. 102 is placed in a tube furnace and connected with a 
bottle containing marble and water from which carbon dioxide 
can be evolved by the addition of hydrochloric acid. As soon as 




Fig. 102. 

the piping is red-hot a slow current of CO^ is allowed to pass 
over the heated charcoal, when carbon monoxide is evolved and 
can be collected over water as shown in the figure. 

ExPT. 124. — Carbon monoxide can be more conveniently pre- 
pared from several compounds of carbon. Thus if crystallised 
oxalic acid (CoH20^) be heated in a flask with strong sulphuric 
acid, a mixture of equal volumes of CO and CO2 are evolved 
(see equation, p. 212) and if these two gases be collected over 
water they can be readily separated by shaking the mixture 
with caustic soda, when half the volume will disappear owing 
to the combination of CO^ with the soda whilst the CO remains 
unabsorbed. The explanation of this mode of preparation is 
that hot sulphuric acid removes the elements of water (HgO) 



LESSON XXI OR CARBONIC OXIDE GAS 223 

from oxalic acid, leaving the residue 0^,0;^ which cannot exist 
alone, and at once splits up into CO and C02- 

ExPT. 125. — A method by which carbon monoxide is obtained 
unmixed with carbon dioxide, is to heat formic acid (CH^O^) or 
sodium formate (CHNa02) with strong sulphuric acid, when 
water (H.^O) is taken up by the acid and CO is evolved as a gas — 

CH.O, = CO 4- Hp. 

Properties. — Carbon monoxide is a colourless, tasteless gas. 
It acts as a strong poison when inhaled even in small quantities, 
and the fatal effect often noticed of breathing the air of rooms 
in which charcoal is burnt in a chauffer, or of gases from lime- 
kilns or brickkilns is due to the presence of this gas. Like all 
other gases, this compound can be liquefied, but to effect this 
a much lower temperature is required than in the case of car- 
bon dioxide, as liquid carbon monoxide boils at- 193°, and 
the gas must, therefore, be cooled below this point before it 
liquefies. The gas is rather lighter than air and cannot be 
collected either by upward displacement, like hydrogen, or 
by downward displacement, like carbon dioxide, but it can be 
collected over water as it is only very slightly soluble in that 
liquid. 

Carbon monoxide burns, when a light is brought to it, with 
a characteristic blue lambent flame with formation of carbon 
dioxide, CO + O = CO2. This serves as a test for this gas. 

ExPT. 126. — Pour some clear lime-water into a bottle filled 
with the gas and observe that, when shaken, the lime-water 
remains clear, then bring a light to the mouth of the bottle, 
observe the flame, and when it has burnt out, shake up the 
lime-water and notice that it becomes milky, and this milkiness 
disappears on the addition of a few drops of hydrochloric acid. 

Carbon monoxide can be separated from other gases by 
bringing it into contact with cuprous chloride, CuCl, when it is 
absorbed. 

Determination of Composition 

The composition of the gas is ascertained by exploding it 
mixed with oxygen (by an electric spark), in this case moisture 
must be present as combination does not occur when the gases 



224 CARBON DIOXIDE OR PART il 

are perfectly dry. One hundred vols, of carbon monoxide yield 

1 GO vols, of carbon dioxide and require 50 vols, of oxygen. 
But, as carbon dioxide contains its own volume of oxygen, 
carbon monoxide must contain half its volume of oxygen, or 

2 vols, weighing 27.8 contain i vol. of oxygen weighing 15.9, 
its formula is, therefore, CO. 

Carbon Dioxide 

Formula CO2. Molecular Weight 43.7. Density 21.85 

Occurrence. — This gas, as we have seen, is formed when- 
ever carbon or any of its compounds burn in excess of air or 
oxygen. It not only exists in the free state in the air, but is 
evolved from the earth, especially near volcanoes, and it accom- 
panies the water in many mineral springs. Moreover, it occurs, 
as has been said, in combination with lime and magnesia, as 
magnesian limestone or dolomite, and with lime as limestone 
coral, chalk, marble, and calcspar. It is not only given off by 
animals in breathing (the expired air from the human lungs 
contains about 4 per cent of this gas), but is evolved in the 
process of decay of animal as well as of vegetable matter, and 
in the fermentation of sugar. Hence, carbon dioxide often 
accumulates at the bottom of caverns, and of old wells and 
mines, as well as in brewers' vats, and also in coal pits, where 
it is known as choke-damp ox after -dainp (damp is the German 
da?npf vapour), formed by the combustion oi fire-da7Jip (see 
p. 230). 

Preparation. — If we want large volumes of carbon dioxide 
we may use the gases coming off from burning coke, or from 
limekilns ; in these, chalk or limestone, calcium carbonate 
(CaCOg) is heated in a draught of air, solid quicklime, CaO, 
remains behind, and CO2 comes off as a gas. 

CaCOg = CaO + CO^. 

But neither of these methods are good for laboratory pur- 
poses, and on the small scale we always prepare this com- 
pound by acting on chalk or marble with an acid — almost any 
acid may be used for this purpose, but hydrochloric acid is the 
best. 

Expt. 127. — Place some lumps of marble in a flask, add 



LESSON XXI 



CARBONIC ACID GAS 



225 



some water, and then pour on to this some hydrochloric acid 
(HCl) through the tube funnel. Efifervescence soon begins, 
owing to the rapid disengagement of the gas, which may be 
collected either over water, or, as it is half again as heavy 
as air, by downward displacement^ as shown in Fig. 103. 

CaCO, + 2 HCl = CaCl, + H^O + CO^. 




The same kind of change occurs if other acids, such as acetic 
acid (vinegar), sulphuric, or nitric acids are added, but then the 
calcium salt corresponding to 
the acid is formed. Sulphuric 
acid is, however, unsuitable 
for the preparation of CO^ from 
marble, because the calcium 
sulphate w^hich is produced, is 
only slightly soluble in water, 
and it forms a coating on the 
marble, thus preventing the 
acid coming in contact with it, 
and thus stopping the evolu- 
tion of gas. Similarly, any other Fig. 103. 
carbonate, such as carbonate of soda (washing soda), or car- 
bonate of potash (pearl-ash), may be used instead of marble, 
•but this latter substance answers best, and is cheapest. 

Properties. — Carbon dioxide is colourless, and has a 
slightly acid taste. It is a heavy gas, being 22 times as dense 
as hydrogen. We can, therefore, syphon it like water, or 
pour it from one vessel to another (see Expt. 2) ; it does not 
support ordinary combustion, and the flame of the taper, as well 
as that of burning phosphorus, are extinguished when plunged 
into the gas. If, however, we heat a bit of potassium in 
a flask filled with the dry carbonic acid gas, the metal is 
seen to take fire and burn. What happens here is that the 
carbon dioxide is decomposed into oxygen, which unites 
with the potassium, to form white fumes of oxide of potas- 
sium and black carbon, which deposits on the side of the glass 
flask. 

Solubility of Carbon Dioxide. — Under the ordinary 
atmospheric pressure carbon dioxide is soluble in water, the 



226 LIQUID CARBON DIOXIDE part il 

maximum which can be dissolved by i gram of water at 
o° is 1.8 vols.; whilst at 20° only 0.9 vols, (or half the 
former amount) dissolves. If, however, the pressure be in- 
creased — that is, if the gas be pumped into water under 
pressure, more of the gas is taken up by the water, and when 
the pressure is removed the gas is evolved. This is well 
seen in the case of ordinary soda water, which contains the 
gas in solution pumped in under a pressure of about 4 atmo- 
spheres. 

ExPT. 128. — If we insert through the cork of a bottle of 
soda water a small screw tap, and attach to the end a piece of 
caoutchouc tubing, we can, by opening the cork, easily collect 
the carbonic acid gas which has been in solution in the water, 
but escapes when the extra pressure is removed. It is easy 
to calculate how much gas will be dissolved in water under 
increased pressure, when we remember that experiment has 
shown that if we double the pressure we double the weight of 
gas dissolved ; and because the volume occupied by any gas 
under varying pressure is inversely proportional to that pres- 
sure, the volM7ne of gas absorbed will remain constant whatever 
be the pressure. Thus if, at the ordinary atmospheric tempera- 
ture and pressure, i gram of water dissolves 2 milligrams of 
CO2, it will dissolve 4 mgm. (or double the amount) when the 
pressure reaches 2 atmospheres, or is also doubled. If soda 
water is made under a pressure of 5 atmospheres, and a 
bottle holds 200 c.c. of water; under the pressure of 760 
mm. the volume of CO2, which this quantity of water can dis- 
solve, amounts to 250 c.c, we see that we can collect 1000 c.c. 
of gas, and this is the difference between the saturating quan- 
tity under the pressure of 5 atmospheres, and that under that 
of I atmosphere. 

As CO2 is formed during the process of fermentation — or 
the conversion of sugar into alcohol by means of a ferment 
such as yeast — w^e find that liquors, such as beer and cham- 
pagne, in which the process of conversion of sugar into alcohol 
is not complete before they are bottled, become saturated with 
carbon dioxide under an increased pressure, and when the 
pressure is diminished by withdrawing the cork, the liquid 
effervesces from escape of the gas. 

Liquid Carbonic Acid Gas. — Like all gases, carbon 



LESSON XXI COMPOSITION OF CO2 227 

dioxide can be condensed to a liquid by the application of 
cold and pressure. The boiling point of the liquid under 
ordinary atmospheric pressure is -78°, cooled still further it 
solidifies to a colourless ice-like mass. If brought under a 
pressure of 35.5 atmospheres at 0°, the gas also liquefies, and 
if the liquid be allowed to escape into the air through a fine 
nozzle attached to the steel cylinder in which it has been 
pumped, part of the liquid at once evaporates, and so much 
heat is thereby absorbed that the rest solidifies. A brass box 
is, for this purpose, attached to the nozzle of the steel cylinder, 
and this, after the liquid has been allowed to enter, is seen to 
be filled with a w^hite snow-like substance w^hich is solid car- 
bonic acid. Liquid carbon dioxide is now a commercial sub- 
stance, sold at IS. per lb., being employed for a variety of 
purposes. The solid is used for the production of very low 
temperatures ; for this purpose the snow-like powder is mixed 
with ether, and the mixture placed in a vacuum ; by this means 
a temperature as low as —100° is attained, and mercury can 
thus be easily frozen. 

Composition of Carbon Dioxide. — How is the compo- 
sition of carbon dioxide ascertained? In the first place, if we 
burn a piece of charcoal in a measured volume of oxygen, and 
take care that neither any oxygen nor any carbon dioxide 
formed in the combustion escape, we shall find that no change 
in the volume of the gas before and after the experiment has 
occurred. This shows that carbon dioxide contains its own 
volume of oxygen, for if it did not do so a change of volume 
must have taken place. Now we know that one volume of 
carbon dioxide weighs 22, when the same volume of hydrogen 
weighs I, or 2 volumes of carbon dioxide wxigh 44; but this 
contains its own bulk (2 volumes) of oxygen, and the 2 volumes 
of oxygen weigh 2 x 16 = 32, so that the w^eight of carbon con- 
tained in 2 volumes of carbon dioxide must be the difference 
(because it contains nothing else), or 44 — 32 = 12; this latter 
number is, however, the atomic weight of carbon, hence carbon 
dioxide is composed of 12 parts by weight of carbon, and 32 
parts by weight of oxygen, or its formula is COo. 

Another method of ascertaining the composition of carbon 
dioxide is to burn a known weight of pure carbon, such as 
diamond or graphite in a current of oxygen gas, collecting and 



228 CARBONIC ACID part ii 

weighing the carbon dioxide formed. This determination has 
been made by several chemists with great care. 

In the first place, the accurately weighed quantity of diamond 
or graphite contained in a small boat of platinum is placed 
inside a porcelain tube, which can be strongly heated in the 
furnace. The oxygen is contained in a gasholder, which is 
connected with a porcelain tube by a series of drying tubes 
by which the gas is rendered perfectly dry and pure before it 
reaches the diamond. At the other end are placed tubes and 
bulbs containing caustic potash to absorb the carbon dioxide 
formed by the combustion of the diamond, whilst other tubes 
contain pumice stone moistened with sulphuric acid to retain 
any moisture which the gas might carry away. Of course 
these tubes are carefully weighed, and then the apparatus put 
together, great care being taken that all the joints are perfectly 
air-tight. After the whole has been filled with dry oxygen, 
the porcelain tube is brought to a red heat, the combustion 
begins and the experiment is allowed to proceed. All the 
carbon dioxide thus generated is absorbed by the caustic potash, 
and, as the oxygen gas is dried both on entering and on leaving 
the apparatus, the gain in weight of the tubes gives the exact 
weight of carbon dioxide produced by the combustion of the 
diamond. As, however, even this purest form of carbon con- 
tains a certain amount of mineral matter which is left behind 
in the platinum boat as ash, it is necessary, in order to obtain 
the exact amount of carbon burnt, to weigh the boat after 
the experiment and to deduct the weight of the ash from that 
of the diamond taken. Another necessary precaution is to 
place some copper oxide in the half of the porcelain tube 
nearest the bulbs and tubes, in order that any carbon monoxide 
(CO) which might be formed and would escape absorption 
by the potash, shall be oxidised to CO^. Many very careful 
experiments made in this way showed that 11.9 parts by weight 
of carbon combine with 31.8 parts of oxygen to form this gas; 
that is, the relation between the number of atoms of carbon 
and oxygen is i to 2, or the formula of carbon dioxide is, as its 
name implies, CO2. 

ExPT. 129. Carbonic Acid and the Carbonates. — If 
a blue litmus paper be plunged into water in which carbon 
dioxide is dissolved, the blue colour is changed to red, showing 



LESSON XXI THE CARBONATES 229 

dioxide is dissolved, the blue colour is changed to red, showing 
the presence of an acid, whilst dry CO^, does not afifect this 
alteration. This shows that the aqueous solution contains an 
acid. Carbonic acid, H^O + CO^ = H^COo, is a dibasic acid 
(see p. 205), but it is very unstable and cannot be obtained in 
the pure state. On the other hand, it forms a series of stable 
salts termed the carbonates. 

ExPT. 130. — When carbon dioxide is passed into a solution 
of caustic soda it is rapidly absorbed, and carbonate of soda is 
formed 

2NaOH + CO2 = Na^COo + H2O. 

Caustic soda. Carbonate of soda. 

If an excess of carbon dioxide be used, or if carbon dioxide 
is passed into a solution of the preceding salt, a neutral solu- 
tion is obtained and another salt, viz. bicarbonate of soda, is 
formed : — 

Na2C03 + CO2 + H2O = 2HNaC03. 

Carbonate of soda. Bicarbonate of soda. 

We thus see that two classes of carbonates exist. One, 
called the normal salt, in which both atoms of hydrogen in 
carbonic acid, H2CO3, are replaced by metal, and the other 
called bicarbonate or the acid salt, in which only one atom is 
thus replaced. 

ExPT. 131. — Let us next pass carbon dioxide through clear 
lime-water, at first a white precipitate of insoluble calcium 
carbonate (CaCOg, see Expt. 4), soluble in hydrochloric acid 
is thrown down, and this serves as a test for carbon dioxide 
After the gas has passed through for a longer time, this white 
precipitate re-dissolves and the liquid becomes clear again 
owing to the formation of a soluble bicarbonate. This com- 
pound is destroyed when the liquid is boiled, CO2 is given off 
and white carbonate of lime (CaCOg) is precipitated (see Expts. 
58 and 59). The carbonates of the alkali metals, potassium 
and sodium, and of ammonium are soluble in water ; the normal 
carbonates of the other metals are insoluble in water. This 
latter fact is seen by adding a solution of carbonate of soda to 



230 MARSH GAS 



a soluble salt of any metal^ such as sulphate of zinc, sulphate of 
copper, or acetate of lead. 



Compounds of Carbon vrith Hydrogens 

These compounds termed the hydrocarbons are extremely 
numerous. They exist as solids, such as paraffin wax ; liquids, 
such as turpentine; and as gases. Of this latter class only 
three, and those most simple, will here be mentioned, viz. 
(i) methane or marsh-gas, CH^, (2) acetylene, C2H2, and 
(3) ethylene or olefiant gas, QH^. 



Methane or Marsh Gas 

Formula CH4. Molecular Weight 15.9. Density 7.95 

This gas is interesting as occurring in coal-pits and causing, 
when mixed with air and the mixture fired, the serious 
explosions which too often occur in coal mines. The pit-men 
term this gas "fire-damp," and the carbonic acid gas which 
results from the explosion "after-damp." Methane is formed 
by the gradual decomposition of organic matter such as leaves, 
and the gas may be seen to arise in bubbles from stagnant pools 
and marshes, and from this the name marsh gas is derived. 
This compound is also a constituent of coal gas, and is not only 
found as fire-damp in coal-pits, but is evolved in such large 
quantities from the oil springs, especially in America, that it is 
carried from the springs for some miles to cities where it is used 
as a source of heat. 

ExPT. 132. Preparation. — If we want to prepare this 
gas in the laboratory we make use of the following reaction : — 

NaC2Ho02 + NaOH = NaaCOg + CH^. 

Sodium acetate and caustic soda give sodium carbonate and 
marsh gas. For this purpose it is best to heat a mixture of 
I part of acetate of soda with 4 parts of soda lime (a mixture 
of caustic soda and lime). The heat required is greater than 
an ordinary glass flask will stand, so it is necessary to use a 



LESSON XXI COMPOSITION OF METHANE 231 

tube of hard glass, closed at one end and fitted with a cork 
and delivery tube at the other. 

Marsh gas cannot be obtained by the direct union of 
carbon and hydrogen, but is formed when a mixture of sul- 
phuretted hydrogen (SH^) and the vapour of carbon disulphide 
(CSo) are passed together over red-hot copper. 

2 H^S + CS2 + 8 Cu = CH^ + 4 Cu^S. 

Properties. — Marsh gas is colourless and tasteless and 
can be condensed to a liquid only by very great pressure, 
the liquid boils at —164^. It has only about half the density 
of air and used to be called light carburetted hydrogen ; it 
is very inflammable, burning with a faint blue flame, forming 
carbon dioxide and water. When mixed with twice its volume 
of oxygen or with ten times its volume of air the mixture 
ignites on application of a flame and a violent explosion occurs ; 
hence the danger of sudden outbreaks of this gas from the 
coal measures, and the necessity of using, in such "fiery 
mines " as they are termed, a safety-lamp (see p. 238) which 
cannot cause the ignition of the combustible gaseous mixture. 

Composition. — Why do we give the formula CH^ to this 
compound.^ The answ^er is given by the following: Let us 
take 2 volumes of this gas and mix it with 6 volumes of 
oxygen in a eudiometer (see p. 92), w^e then find that after 
passing an electric spark through the mixture the 8 vol- 
umes have been reduced to 4 volumes. Next let us absorb 

CH^ + 3O2 = 2H2O + CO2 + O2. 

2 vols. 6 vols. 2 vols. 2 vols. 

the carbon dioxide formed by the above combustion by 
means of caustic soda; we find that 2 volumes of oxygen 
gas remain. Hence 4 volumes of oxygen were needed to 
burn the carbon and hydrogen contained in 2 volumes of 
methane, of which 2 have gone to burn the carbon to carbon 
dioxide and 2 to burn the hydrogen to water. Therefore 2 
volumes of methane must contain 4 volumes of hydrogen 
weighing 4, and the weight of carbon (viz. 11.9 parts) which 
is contained in two volumes of CO2, and since its density is 
7.95, we conclude that CH^ is its molecular formula. 



232 ACETYLENE part ii 

Acetylene 

Formula C2H2. Molecular Weight 25.8. Density 12.9 

This gas is of interest because it is formed by the direct 
union of its elements, whilst the other two compounds are 
produced indirectly. Carbon and hydrogen can only be made 
to unite at very high temperatures, such as that of the electric 
arc. If the two carbon poles of a powerful battery or dynamo are 
inclosed in a vessel filled with hydrogen, and a strong electric 
current passed from pole to pole, acetylene is formed. It is a 
colourless gas possessing a peculiar and disagreeable smell, 
it burns with a smoky luminous flame when a light is brought 
in contact with it. Acetylene is also formed in cases where 
the combustion is incomplete, as w^hen, for instance, the Bunsen 
burner " burns down," that is, when the coal-gas burns from 
the jet at the bottom of the tube instead of the mixture of 
air and gas burning at the top (see p. 236). 

ExPT. 133. — To show the formation of this gas, place a 
glass tube 12 inches long over the Bunsen w^hen burning down, 
and invert over this tube a large flask, the inner surface of 
which has been moistened with an ammoniacal solution of 
cuprous chloride.* The blue colour of this solution will in a 
few moments be changed to a deep red, owing to the formation 
of a red insoluble compound of acet3'lene with cuprous oxide. 
This is an excellent test for the presence of this compound. 
Acetylene combines directly with hydrogen to form the next 
compound, ethylene, C2H2 + Hg — CgH^. 

Ethylene or defiant Gas 

Formula C2H4. Molecular Weight 27.8. Density 13.9 

This gas is one of the chief constituents of the product of 
the destructive distillation of coal (see coal-gas, p. 234), and it is 
chiefly to its presence that coal-gas owes its luminous qualities. 

Preparation. — In order to prepare ethylene in a pure 
state, I part of alcohol or spirits of wine (CgHgO) is heated 

* Cuprous Chloride, CU2CI2, is prepared by the action of metallic copper 
on a boiling solution of cupric chloride, CUCI2, and hydrochloric acid. 



LESSON XXI ETHYLENE 233 

in a flask with 5 or 6 parts of strong sulphuric acid, this 
mixture being first made by carefully pouring the alcohol into 
the sulphuric acid in a thin stream, whilst the mixture is 
constantly stirred. In this experiment the flask should 
previously be half-filled with dry sand and the mixture poured 
on to it as otherwise the liquid is apt to froth over. On 
heating the flask, care being taken that the sand at the bottom 
of the flask is wxtted with the mixture, ethylene gas is quickly 
evolved and may be collected as usual over water. The 
reaction that here occurs is similar to that by which carbon 
monoxide is prepared from formic acid (p. 223), the sulphuric 
acid removes the element of water (H2O) from the alcohol and 
C2H4 is evolved. 

Properties. — Ethylene is a colourless gas having a sweet- 
ish taste; when cooled to —110° it condenses to a colourless 
liquid which at —160^ solidifies. It burns with a brightly 
luminous flame when a light is applied. To show this best, 
a cylinder filled with the gas is opened, a lighted taper applied 
to the mouth, and water quickly poured in, when a large smoky 
flame issues with formation of carbon dioxide and water. 

When mixed with three times its volume of oxygen and 
fired, the mixture detonates very strongly. Care must be 
taken in making this experiment. The above mixture may 
be made in an ordinary glass flask, and the flask then corked 
and wrapped up in a strong towel, care being taken to leave 
the mouth of the flask exposed. On removing the cork and 
bringing a lighted taper to the mouth a violent explosion 
occurs and the flask is shattered to pieces. Unless the above 
precautions are taken the glass of the flask is thrown about 
and serious results w^ould follow. 

Two volumes of olefiant gas, C2H4, require for their com- 
plete combustion 6 volumes of oxygen ; 4 volumes for Cn, 
forming 4 volumes of COg, and 2 for the hydrogen, forming 4 
volumes of water as shown in the equation. 

C2H4 + 3 O2 = 2 CO2 + 2 H2O. 

2 vols. 6 vols. 4 vols. 

Ethylene derives its old name, olefiant, or oil-making, gas 
from the fact that it unites with its own volume of chlorine to 
form an oily liquid, C2H^Cl2, ethylene dichloride. 



234 COAL GAS part ii 



Coal Gas and Flame 

This substance, manufactured on so large a scale for 
illuminating and heating purposes, is not a chemical compound 
but a mixture of many compounds. If bituminous coal be 
heated to redness in a closed retort three distinct kinds of pro- 
ducts are given off: (i) gas tar; (2) ammonia liquor or gas 
water; and (3) gas. Tar is a mixture of a great variety of 
chemical compounds, many of which now yield very valuable 
products, such as the aniline colours, besides scents and useful 
medicines. The ammonia liquor is the chief source of the 
ammoniacal salts ; it is derived from the nitrogen which the 
coal contains. The gas contains several ingredients, some of 
which are harmful, and must be removed, whilst others are 
useful and must be retained. Amongst the former are carbon 
dioxide (CO2), sulphuretted hydrogen (SHo), and the vapour of 
carbon disulphide (CS2), and these impurities are more or less 
completely removed from the gas before it is sent out from the 
gasworks for consumption. The useful ingredients may be 
divided into two classes: (i) those which act as illuminating 
agents ; and (2) those which burn and give out heat, but do not 
give off light when burning. The first class consists chiefly 
of ethylene, mixed with small quantities of acetylene, CgHg, and 
the vapours and other hydrocarbons such as propylene (CoH^.) 
and benzene (QH^j) ; in the second class are contained 
hydrogen, carbonic oxide, and methane or marsh gas (CH_j), 
these serve as diluents to dilute the ethylene, which by itself 
burns with too smoky a flame to be available for ordinary 
purposes. Various kinds of coal yield gas which differs both 
in composition and in illuminating power ; moreover the heat 
to which the coal is subjected greatly affects the properties 
and composition of the gas evolved. 

The degree of luminosity of coal-gas is ascertained by com- 
parison with the light given out by a standard candle. Of 
course the size of the flame, or the rate at which the gas is 
burning, must be taken into account. For this purpose the 
standard used is a flame consuming 5 cubit feet of gas per 
hour. When a coal-gas is said to be equal to 17.5 candles, this 
means that a gas flame burning at the above rate gives off an 



LESSON XXI 



NATURE OF FLAME 



235 



amount of light 17.5 times as great as that of one standard 
candle. 

Water-gas is a mixture of hydrogen and carbon monoxide 
obtained by passing steam over red-hot coal or coke^ when 
the following decomposition occurs : — 

C + H.O = CO + H^. 

This gas burns with a non-luminous but very hot flame, 
and it is thei;efore used for steel-making and for other purposes, 
where a high temperature is needed. For illuminating purposes 
the flame is allowed to impinge on a comb of magnesia, wdiich 
becomes white hot, and then emits a powerful light. 



The Structure of Flame 

Flame is gas in a highly-heated or incandescent state. If 
flame passes quickly through an inflammable mixture, such as 
oxygen and hydrogen, or coal-gas and air, or if a rapid com- 
bustion of a solid or liquid body, such as gunpowder or nitro- 
glycerine, takes place, the sudden expansion gives rise to an 
explosion. If, however, the inflammable gases come slowly into 
contact with the air, as in the case of a jet of coal-gas, or of a 
lighted lamp or candle, a steady flame is seen. A candle flame 
consists of three parts or zones — (i) the 
dark central zone consisting of the supply 
of combustible and unburnt gas surround- 
ing the wick ; (2) the luminous zones ; (3) 
the non-luminous exterior zone. 

ExPT. 134. — By bringing the end of 
the bent tube (Fig. 104) into the interior 
zone the unburnt gases may easily be with- 
drawn from the centre of the flame, and 
lighted at the longer end of the syphon- 
tube. In the luminous zone the combustion 
is incomplete, and carbon in the form of 
soot is separated out. This is easily shown 
by holding a card or sheet of stiff" paper 
horizontally in the candle flame for a few moments, so as 
not to burn the paper, when on withdrawing the card a black 




236 



THE BLOWPIPE FLAME 



PART II 



ring of soot is seen to be deposited, whilst the centre remains 
white. The luminosity of a flame depends on the presence 
of solid particles ; where the combustion is complete, that is 
where all the carbon is at once burnt up to carbon dioxide as 
in the outer zone (3) of the candle flame, no luminosity occurs, 
but if we bring a thin piece of platinum wire into this part of 
the flame, we see that the wire becomes white hot, and gives 
off light. The effect of complete combustion rendering the 
flame non-luminous is seen in the Bunsen burner, now gen- 
erally used in laboratories. Its construction is seen in Fig. 
105 ; the coal-gas issues from a small burner (a) at the foot 
of the lamp, and passes up unburnt to the top of the tube 
(c), where it mixes with air drawn up the tube through the 
holes {dd) ; on bringing a light to the top of the 
tube, the mixture of air and gas burns with a 
blue, perfectly smokeless flame : but if the holes 
(dd) are closed by the fingers, the gas burns 
with the ordinary luminous flame. This non- 
luminous flame, like a candle flame, is hollow. 
To show this, several experiments can be made. 
ExPT. 135. — Take a thin platinum wire, and 
hold it horizontally in the Bunsen flame. It will 
then be noticed that the wire becomes brightly 
luminous in two points at each side of the flame, 
whilst the part of the wire in the centre of the 
flame does not glow. Again, thrust the head of a lucifer match 
quickly through a Bunsen flame to the centre when it will be 
seen that the tip may be held for some time in the central 
zone without taking fire, whilst 
the wood will be charred where it 
comes in contact with the heated 
outer mantel. 

ExPT. 136. The Blowpipe "'"* ^"^^^ 

Flame. — Place a hollow tube, 
having the upper end flattened 
and sloped, inside the tube of the ^^^' ^°^* 

Bunsen, so that the air is prevented from entering the holes 
{dd Fig. 105), and then urge the luminous flame with a mouth 
blowpipe. It will be seen that, like the candle flame, that of 
the blowpipe consists of two parts — (i) the oxidising or outer 




Fig. 105. 




LESSON XXI THE DAVY LAMP 237 

part of the flame (a) ; and (2) the reducing or inner part (d). 
In the first there is an excess of oxygen, and the flame is non- 
luminous ; in the second there is an excess of carbon, and the 
flame is luminous. To show the different effects which these 
two parts of the flame exert, let us make a colourless borax 
bead on a loop of platinum wire, by heating a little powdered 
borax on the loop until it fuses to a colourless glass. Then 
dip this bead into a solution of ferric chloride (Fe^Cl^) ; on 
heating the bead in the outer flame it will be seen that the 
borax-glass becomes coloured yellow ; next heat it carefully 
in the inner or reducing flame, and it will be seen that the 
colour of the bead is changed to green. This is due to the 
reduction of the ferric oxide (FcoOo) (formed in the outer 
flame) to ferrous oxide (FeO) in the inner flame. 

The Davy Lamp. — In order that a gas shall become ignited, 
the temperature of the gas must be raised to a certain point, 
below this point the gas will not inflame. Different gases and 
vapours ignite at very different temperatures. 

ExPT. 137. — To show this, try to light a jet of coal-gas 
with a red-hot splinter of wood. You will not succeed in doing 
this, but blow the red-hot splinter into a flame, and the gas 
at once ignites. Then pour a few drops of bisulphide of carbon 
(CS2) into a saucer and try the same experiment. The red- 
hot splinter of wood immediately ignites the vapour, which 
then burns with a lambent-blue flame with evolution of 
SO, + CO,. 

A candle flame may be so cooled down as to extinguish it. 
For this purpose make a coil of cold copper wire, and place it 
over the flame ; the flame will at once go out. 
Then heat the coil and repeat the experiment, 
when the flame will continue to burn. 

It is upon this principle that the ^' Davy " 
or safety-lamp for coal-miners is con- 
structed. To understand this, let us make 
a simple experiment. Take a square piece 
of wire gauze, containing about 700 meshes 
to the square inch. Hold this over a jet 
of gas, and light the gas on the top of ^^s- 107. 

the gauze ; then carefully raise the gauze, and observe that 
whilst the flame continues to burn on the upper surface, the 




238 



SUMMARY 



PART II 



flame is seen not to pass through the gauze, and no flame 
exists below it. The metal wires here so quickly conduct 
away the heat of the flame that the temperature of the in- 
flammable gas below the gauze never reaches the point of 
ignition, and the gauze may be moved upwards so that at last 
the flame is extinguished, although the gas continues 
to escape from the jet. Now imagine this gauze 
wrapped round a lighted candle, both above, below, 
and at the sides. The air can get through the 
meshes, and the products of the combustion can 
escape, but if you approach the outer side of the 
gauze with an unlighted jet of gas, you will not be 
able to light the jet. This is the "Davy Lamp." 
Its construction is seen in Fig. 108. A gauze 
cover is screwed on to the rim of an oil lamp, and 
when the lamp is lighted, and the cover screwed 
on, it may be plunged into a large beaker, at the 
bottom of which a small quantity of ether has been 
poured, and yet the inflammable vapour of the ether 
will not become ignited. Then remove the safety- 
lamp, and throw into the beaker a bit of burning paper, and 
the whole vessel will at once be filled with flame. This experi- 
ment illustrates the use of this lamp in coal-pits where inflam- 
mable mixtures of " fire-damp " and air are liable to occur. 




Fig. 108. 



What we have learnt 

In our twenty-first Lesson we have studied the various allotropic forms 
of carbon, viz. diamond, graphite, and charcoal in its various forms. 
Carbon monoxide is prepared by the action of strong sulphuric acid on 
formic acid or oxalic acid, in the latter case the CO is mixed with an equal 
volume of CO2. Carbon monoxide is also produced when charcoal is 
burnt in an insufficient supply of air, or when carbon dioxide is passed 
over red-hot charcoal. CO is a colourless poisonous gas, which burns in 
the air with a lambent blue flame forming CO2. Its composition may be 
determined by exploding it with oxygen in a eudiometer, when it is found 
that 2 volumes of CO (density 13.9) combine with i volume of oxygen to 
form 2 volumes of CO2. 

Carbon dioxide is prepared by the action of hydrochloric acid on marble^ 
or it is produced whenever carbon or carbon compounds burn in excess of 
air. It is sUghtly soluble in water, the amount dissolved being proportional 
to the pressure at which solution is effected. It exists dissolved under pres- 



LESSON XXI EXERCISES 239 

sure in all effervescent beverages, and, although sometimes introduced 
artificially, it is produced naturally in other cases by a process of fermenta- 
tion. CO2 is a very heavy gas, and may be collected by downward dis- 
placement. Its composition is ascertained by the same method as was 
used for sulphur dioxide. 

CO.2 is recognised by its power of rendering lime-water turbid. It is the 
anhydride of carbonic acid, H2CO3, which, like sulphurous and sulphuric 
acids, is dibasic, and forms therefore two series of salts, viz. M'2C03 and 
M'HCOs. 

The hydrocarbons marsh gas, CH4, acetylene, C2H2, and ethylene, 
C2H4, have been shortly considered, as also has coal gas which is a mixture 
of various gases obtained by the destructive distillation of coal. Finally, 
we have studied the structure of flame, and its chemical characters. We 
have seen how a knowledge of the characters of flame, and the conditions 
under which it is produced, led Sir Humphry Davy to devise the safety- 
lamp for miners. 

Exercises on Lesson XXI 

1. By what three methods could you prepare carbon monoxide ? 

2. What are the properties of carbon monoxide and dioxide, and how 
may their composition be determined ? 

3. How would you distinguish carbon monoxide from hydrogen and 
from marsh gas ? 

4. How would you prepare marsh gas and ethylene ? 

5. Why is it that miners often perish in a coal mine after an explosion 
of fire-damp, although unhurt by the actual explosion ? 

6. Describe the eudiometric analysis of marsh gas, giving the volume 
ratios. 

7. Explain the burning of a candle. 

8. By what experiments would you illustrate the principle of the Davy 
safety-lamp ? 

9. What volume of CO2 at NTP can be obtained from 185 grams of 
marble (i) by decomposing it by heat; (2) by acting on it by dilute nitric 
acid? 

10. By what tests would you recognise that a certain gas is a compound 
of carbon and hydrogen ? 

11. What is water gas? Give an equation showing its formation. What 
are the products of its combustion in (i) air, (2) oxygen ? 

12. How is bi-carbonate of soda prepared ? What happens when CO2 
is passed through lime-water (i) in small quantity, (2) in large quantity, 
and what happens in the last case when the liquid is boiled ? 



242 



APPENDIX 







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INDEX 



Acetylene, 232 

Acids, 12, 81 

Acids and salts, nomenclature of, 

82 
Acid salts, 205 
Air, 8, 131 

" a mixture, 134 

" dissolved in water, 121 

" pressure of, 131 
Alkali, 15 
Ammonia, 162 

" freezing machine, 165 

Ammonium salts, 168 
Analysis and synthesis, 13 
Aqua regia, 178 
Atomic theory, 30 
" weight, 31 
Avogadro's law, 59 

Balance, the, 39 
Barometer, 52, 132 
Basic oxides, 81 
Bleaching powder, 185 
Blowpipe flame, 236 
Bluestone, 23 
Boiling point, 108 
Boyle's law, 52 
Burette, 44 

Calculations of weights of mate- 
rials, 34 
Calorie, 108 
Candle, experiments with, 17 



Carbon, 216 

" dioxide, 224 

" monoxide, 211 
Carbonates, 228 
Carbonic acid, 228 

" in air, 139 

Carbonic oxide, 221 
Carre's machine, no 
Carrier of oxygen, 208 
Charcoal, 219 
Charles's law, 250 
Chlorides, 184 

" separation of, 185 

Chlorine, 170 
Clark's process, 124 
Coal, 220 
Coal-gas, 234 
Combination, 12 
Combining volumes of gases, 59 

" weights, 28 

Combustible body, 79 
Compounds, 21, 23 
Constancy of composition, 27 
Corrosive sublimate, 23 
Critical point of liquefaction, 64 
Cryophorus, no 
CrystaUisation, 118 

" water of, 119 

Dalton's atomic theory, 30 
Dalton's law, 50 
Decomposition, 12 
Deliquescence, 119 



243 



244 



INDEX 



Dew-point, 142 
Diamond, 218 
Diffusion of gases, 60 
Dissolved air, 121 
Distillation, 125 * 

Distribution of elements, 22 
Dulong and Petit's law, 112 
^ Dyads, 82 

Earth's crust, composition of, 

23 
Efflorescence, 119 
Electrodes, 16 
Electrolysis of water, 15, 95 
Electrolytic gas, 96 
Elements and compounds, 21 
Equations, 25 
Ethylene, 232 
Eudiometer, 92 
' Evaporation, 105 
Expansion of gases, solids and 
liquids, 44 

Flame, 234 

Formulae, 25, 34 

" calculation of, 37 

Freezing machines — 

" Ammonia, 165 

" Carre's, no 

Freezing mixtures, in 

Gases, calculations, 50-59 

" experiments with, 1-3 
Gay-Lussac's law, 55 
Graham's law of diffusion, 63 
Gramme, 39 
Graphite, 219 
Gravitation, 39 

Hard water, 123 
Henry's law, 120 
Hope's experiment, 102 
Hydrochloric acid, 175 
Hydrogen, 67 

" dioxide, 99 

" disulphide, 199 

" peroxide, 99 

" sulphide, 192 



Hydrosulphuric acid, 192 
Hydroxides, 81 

Indestructibility of matter, 17 
Indicator, 83 

Kilo, 38 

Kipp's apparatus, 193 

Lamp-black, 220 

Latent heat, 106 

Laughing gas, 153 

Lavoisier's experiments on air, 9 

" " on water, 99 

Law of diffusion, 63 
Leaden-chamber crystals, 209 
Lead-tree, 24 
Length, standards of, 38 
Lime-light, 79 
Liquefaction of gases, 63 
Litmus an indicator, 83 
Litre, 39 
" of hydrogen, weight of, 56 

Marsh gas, 230 
Mass, 39 

Matter indestructible, 17 
Maximum density of water, 102 
Mechanical mixture, 16 
Mercury solidified, 6 
Metals and non-metals, 22 
Methane, 230 
Metre, 38 
Molecule, 31 
Monads, 82 

Multiple proportions, combination 
in, 29 

Nascent state, 175 
Neutralisation, 83 
Nitrates, 151 
Nitric acid, 149 

" anhydride, 160 
" oxide, 153 
Nitrogen, 9, 129 

" five oxides of, 30 

" monoxide, 153 

" dioxide, 156 



INDEX 



245 



Nitrogen, trioxide, 159 

" tetroxide, 160 

" pentoxide, 160 

" peroxide, 160 
Nitrous acid, 159 

" oxide, 153 
Normal salts, 205 

Occurrence of the elements, 22 
Oil of vitriol, 207 
Olefiant gas, 232 
Oxides, 81 

Oxyhydrogen light, 79 
Oxygen, 74 

" from air, 11, 77 
Ozone, 85 

" rate of diffusion of, 89 

Percentage composition, calcula- 
tion of, 34 
Permanent gases, 64 

" hardness of water, 124 

Peroxides, 83 
Precipitation, 117 
Phosphorus burnt in air, 8, 129 
Physical changes, 6 

Reduction to unity, 35 
Rider, use of, 41 

Salts, 81 

Size of molecules, 31 
Sodium and water, 14, 70 
Soft water, 123 
Solidification of gases, 65 
Solids, liquids, and gases, i 
Solubility of gases, 120 

" of salts, 116 
Specific gravity, 113 

heat, 112 
Steam, action of, on iron, 70 
" composition of, 94 



Steam, invisible, 5 
Storage of oxygen, 'jj 
Sulphates, 213 
Sulphides, 197 
Sulphites, 205 
Sulphur, 187 

" and iron, 16 

" dioxide, 200 

" trioxide, 205 
Sulphuretted hydrogen, 192 
Sulphuric acid, 207 

" anhydride, 205 

Sulphurous acid, 204 

" acid gas, 200 

" anhydride, 200 

Supporter of combustion, 11, 79 
Symbols, 24, 32 
Synthesis and analysis, 13 

Temporary hardness, 123 
Thermal Unit, 108 
Thermometer, graduation of, 45 
Thermometric scales, 47 
Three states of matter, 44 
Torricellian vacuum, 52 

Vapour tension, 105 
Ventilation, 145 
Vitriolic acid, 207 
Volumes, measurement of, 43 

Water as a solvent, 116 

" composition of, 91-99 

" hard and soft, 123 

" heat relations, 102 

" introductory, 14 

" maximum density of, 102 

" natural, 121 

" of crystallisation, 119 

" purification of, 125 

Water-gas, 235 

Weight, standard of, 39 



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